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THE  LIBRARY 

OF 

THE  UNIVERSITY 

OF  CALIFORNIA 

DAVIS 

GIFT  OF 


ROBERT  I.   TEMEY 


r  f  ^ '  ^    '3 


QUANTITATIVE  ANALYSIS 


:''^i^>M^' 


THE  MACMILLAN  COMPANY 

NEW  YORK   •    BOSTON   •   CHICAGO 
DALLAS  •   ATLANTA   •    SAN  FRANCISCO 

MACMILLAN  &  CO.,  Limited 

LONDON   •    BOMBAY   •   CALCUTTA 
MELBOURNE 

THE  MACMILLAN  CO.  OP  CANADA,  Ltd. 

TORONTO 


A  COURSE  OF  INSTRUCTION   IN 

QUANTITATIVE 
CHEMICAL    ANALYSIS 

FOR  BEGINNING  STUDENTS 

with  explanatory  notes,  questions  and 
analytical  problems 

by' 
GEORGE  McPHAIL  SMITH 

PROFESSOR  OP  CHEMISTRY  IN  THE  UNIVERSITY  OF  WASHINGTON 


REVISED  EDITION 


THE  MACMILLAN  COMPANY 
1922 

All  rights  reserved 


Copyright,  1919  and  1921, 
By  the  MACMILLAN  COMPANY. 


Set  up  and  electrotyped.    Published  July,  1919. 
Revised  edition,  Published  October,  192 1. 


PRINTED  IN  THE  UNITED  STATES  OF  AMERICA 


PREFACE 

This  book  is  intended  for  use  with  classes  consisting  of  students 
who  have  completed  a  substantial  year's  work  in  Elementary  Chemis- 
try and  Qualitative  Analysis,  and  who  are  beginning  the  study  of 
Quantitative  Analysis.  On  the  laboratory  side  its  main  object  is  to 
furnish  the  student  with  directions  sufficiently  detailed  to  enable 
him  to  do  successful  work  with  a  minimum  of  personal  oversight  from 
the  instructor.  The  instructor  is  thereby  placed  in  a  position,  in  the 
laboratory,  as  well  as  in  the  classroom,  to  exert  his  efforts  more  es- 
pecially towards  the  development  in  his  students  of  theoretical  knowl- 
edge and  the  power  to  think. 

For  the  best  results,  the  laboratory  work  should  of  course  be  ac- 
companied by  lectures  and  recitations ;  and,  with  the  latter  in  mind, 
it  has  seemed  desirable  to  include  the  questions  of  Part  IV,  and  the 
stoichiometrical  problems  of  Part  V.  *  The  questions  of  Part  IV  are 
for  the  most  part  answered  in  the  notes  or  elsewhere  in  the  book; 
but  it  has  been  the  writer's  experience  that  beginners  react  more 
favorably  to  concrete  questions  assigned  in  advance  for  study,  than 
to  the  same  questions  when  put  to  them  for  the  first  time  just  after 
they  are  supposed  to  have  mastered  the  details  and  principles  of  a 
specific  analytical  process.  The  problems  of  Part  V  are  such  as  are 
regularly  met  with  in  analytical  work,  and  their  conscientious  study 
will  give  the  student  an  insight  into  the  principles  of  a  wide  variety  of 
processes;  the  answers  to  the  problems  have  been  intentionally 
omitted.  It  has  been  the  writer's  practice  to  require,  as  a  written 
exercise  to  be  handed  in  at  the  beginning  of  a  recitation,  the  solution 
of  a  definite  number  of  problems  each  week  throughout  the  course ; 
these  to  be  graded  and  returned  at  the  end  of  the  following  recita- 
tion. 

Part  I  comprises  a  somewhat  detailed  discussion  of  those  principles, 
both  of  theory  and  practice,  which  should  be  constantly  kept  in  mind 
by  the  worker  in  analytical  chemistry.  It  is  realized  that  the  student 
can  assimilate  this  material  only  gradually,  as  his  progress  advances ; 


vi  PREFACE 

but  the  writer  feels  strongly  that  such  a  discussion  should  be  avail- 
able, especially  in  an  introductory  manual. 

The  gravimetric  and  volumetric  processes  detailed  in  Parts  II  and 
III  are  such  as  have  for  several  years  been  included  in  the  elementary 
courses  of  quantitative  analysis  at  the  University  of  Illinois.  On 
the  basis  of  a  good  deal  of  experience,  they  are  regarded  as  processes 
which  are  well  adapted  to  the  progressive  instruction  of  beginners, 
and  are  beHeved  to  constitute  an  excellent  foundation  for  the  con- 
tinued study  of  analytical  chemistry.  Also,  it  is  beHeved,  they  are 
sufficiently  varied  to  meet  the  needs  of  students  who  will  not  extend 
their  study  beyond  the  period  of  an  introductory  course. 

In  addition  to  the  help  derived  from  other  books  and  from  journal 
articles,  in  the  preparation  of  this  manual,  the  writer  wishes  to  ac- 
knowledge his  indebtedness  to  the  following  works  on  analytical 
chemistry:  W.  C.  Blasdale's  Principles  of  Qimntitative  Analysis; 
A.  Fischer's  Elektroanalytische  Schnellmethoden;  W.  F.  Hillebrand's 
The  Analysis  of  Silicate  and  Carbonate  Rocks;  E.  G.  Mahin's  Quanti- 
tative Analysis;  J.  W.  Mellor's  Treatise  on  Quantitative  Inorganic 
Analysis;  H.  P.  Talbot's  Quantitative  Chemical  Analysis;  and  F.  P. 
Treadwell's  Lehrbuch  der  analytischen  Chemie. 

G.  McP.  Smith. 

Seattle,  Washington 
1921 


CONTENTS 

PART   I 
INTRODUCTION 

PAGE 

A.  Gravimetric  and  Volumetric  Analysis i 

B.  General  Remarks  Concerning  Quantitative  Work      .        .        3 

Physical  Conditions;  Reagents  and  Glassware;  Utilization  of 
Time ;  Accuracy ;  Laboratory  Records. 

C.  The  Operations  of  Analytical  Chemistry 

I.  Weighing 7 

Standards  of  Mass;  The  Balance;  The  Use  and  Care 
of  the  Analytical  Balance;  Determination  of  the  Rest- 
point;  Methods  of  Weighing;  The  Calibration  of  a 
Set  of  Weights;  Errors  Due  to  Inequalities  in  Length 
in  the  Beam  Arms ;  Errors  Due  to  Atmospheric  Buoy- 
ancy; Summary. 

II.  Precipitation 21 

Qualities  Desirable  in  Precipitates  Which  Are  to  be  Used 
in  Gravimetric  Determinations;  Colloidal  and  Fine- 
grained Precipitates ;  The  Contamination  of  Precipitates ; 
The  Theory  of  Precipitation. 

III.  Filtration  and  the  Washing  of  Precipitates  .        .        .31 

The  Selection  and  Use  of  Paper  Filters ;  Wash  Bottles ; 
Gooch's  Filtration  Crucible;  The  Theory  of  Washing 
Precipitates. 

rV.  The  Drying  and  Ignition  of  Precipitates      ....      38 

Drying  Ovens ;  Desiccators ;  Crucibles. 
V.  The  Evaporation  of  Liquids 

VI.  The  Volumetric  Measurement  of  Liquids     . 

Volumetric  Apparatus;  Sources  of  Error  in  the  Use  of 
Volumetric  Apparatus;  The  Calibration  of  Volumetric 
Apparatus. 

D.  The  Preparation  of  Samples  for  Analysis   . 

vii 


SO 


viii  CONTENTS 

PART  II 
GRAVIMETRIC   ANALYSIS 

FAGB 

Exercises  with  the  Balance 53 

The  Determination  of  Chlorine  in  a  Soluble  Chloride  .        .      55 

The  Determination  of  Iron  and  of  Sulphur  in  a  Soluble  Sul- 
.  PHATE  OP  Iron 60 

The  Determination  oa  Sulphur  in  a  Sulphide  Ore  ...  66 
The  Determination  ob  Potash  in  Soluble  Salts       ...  68 
The  Determination  op  Carbon  Dioxide  in  Limestone       .        .71 
The  Determination  of  Calcium  and  Magnesium  Oxides  in  Lime- 
stone   75 

The  Determination  of  Phosphoric  Anhydride  in  Phosphate 

Rock  81 

The  Determination  of  Silica  in  a  Refractory  Silicate  .        .  85 

The  Electrolytic  Determination  of  Copper       ....  89 

PART  III 
VOLUMETRIC  ANALYSIS 

General  Discussion »       .        »     99 

Fundamental  Principles;  The  Fundamental  Reactions  of  Volu- 
metric Analysis;  Determination  of  the  End-Point;  General 
Theory  of  Indicators;  Volumetric  vs.  Gravimetric  Methods; 
General  Remarks. 

A.  Neutralization  Methods:  Acidimetry  and  Alkalimetry      .    107 
Fundamental   Principles;    Methyl   Orange   Solution;    Phenol- 
phthalein  Solution ;   Standard  Acid  Solutions ;  Standard  Alkali 
Solutions. 

The   Preparation  and   Standardization   of   Approximately 
Half-normal  Hydrochloric  Acid  and  Sodium  Hydroxide      .    112 

The  Determination  of  the  Alkaline  Value  of  Soda     .        .116 
The   Determination  of  the  Available  Hydrogen-ion  in  an 
Acid 118 


CONTENTS  ix 

PAGB 

The  Determination  of  Protein  Nitrogen  by  the  Kjeldahl 

Method 120 

B.  Methods  of  Oxidation  and  Reduction 124 

Standard  Solutions ;  Indicators. 

1.  DiCHROMATE  PROCESSES 1 25 

Fundamental  Principles. 
The  Preparation  and  Standardization  of  Approximately 
Tenth-normal  Bichromate  and  Ferrous  Iron  Solu- 
tions .        .        .        .       ■ 125 

The  Determination  of  Iron  in  Siderite    .        .       .        .129 

2.  Permanganate  Processes 130 

Fundamental  Principles. 
The  Preparation  and  Standardization  of  an  Approxi- 
mately Tenth-normal  Solution  of  Potassium  Per- 


manganate         

The  Determination  op  Iron  in  Hematite  . 

The  Determination  op  Calcium  in  Limestone  . 

The  Determination  op  the  Mn02- Value  of  Pyrolusite 

The  Determination  op  Phosphorus  in  Steel     . 

The  Determination  op  Manganese  in  an  Ore  . 

3.  Iodometric  Processes 

Fundamental    Considerations;     Determination    of    the 
End-Point ;  Preparation  of  Starch  Solution. 

The  Preparation  and  Standardization  of  Approxiaiately 
Tenth-normal  Solutions  of  Iodine  and  Sodium  Tmo- 
sulphate 

The  Determination  of  Antimony  in  Stibnite 

The  Determination  of  Chromium  in  Chromite 

The  Determination  of  Lead  in  an  Ore 

The  Determination  of  Copper  in  an  Ore 


131 
133 
137 
138 
140 
144 
148 


151 

157 
159 
162 


C.  Precipitation  Methods 165 

General  Discussion. 

The  Preparation  and  Standardization  of  Approximately 
Tenth-normal  Solutions  of  Silver  Nitrate  and  Am- 
monium Thiocyanate 166 

The  Determination  of  Chlorine  in  a  Soluble  Chloride     168 


X  CONTENTS 

PART  IV 

PAGE 

Questions ^^^ 

PART  V 

ANALYTICAL  PROBLEMS 

Preliminary  Discussion:  The  Solution  of  Typical  Problems  .    185 

Problems  j  2 

APPENDIX 

Preparation  of  the  Reagents 20^ 

Sulphuric  Acid-Dichromate  Cleaning  Solution  .        .        .        .208 
Analytical  Samples  for  the  Use  of  Students    .        .        .        .208 

Apparatus  in  the  Student's  Desk 200 

Logarithms 210 

Antilogarithms 212 

Index 

215 

International  Atomic  Weights,  1921     .        .        .        Back  Cover  Sheet 


PART  I 

INTRODUCTION 
A.   GRAVIMETRIC  AND  VOLUMETRIC  ANALYSIS 

Quantitative  analysis  has  for  its  object  the  determination  of 
the  quantities  of  the  elements  or  compounds  which  are  present 
in  particular  samples  of  material.  The  results  are  usually  ex- 
pressed in  terms  of  percentage,  ordinarily  by  weight ;  but  some- 
times, as  in  the  analysis  of  gases,  by  volume. 

The  procedure  to  be  employed  in  a  specific  case  will  often  de- 
pend upon  the  quaHtative  composition  of  the  sample.  A 
quaHtative  analysis,  therefore,  should  precede  a  quantitative, 
unless  the  composition  of  the  sample  is  sufficiently  well  known. 

In  the  performance  of  quantitative  determinations  there  are 
two  principal  methods  of  procedure,  according  to  which  the  sub- 
ject is  subdivided  into  gravimetric  and  volumetric  analysis.  In 
addition,  there  are  gasometric  methods,  and  various  physical 
methods,  of  analysis ;  but  these  will  not  be  described  in  this  book. 

In  a  gravimetric  analysis,  a  weighed  sample  is  taken,  and  the 
substances  to  be  determined  are  separated,  one  after  another, 
either  in  the  free  state,  or  in  the  form  of  suitable  compounds. 
Each  final  product  is  weighed,  and,  from  its  weight,  the  weight, 
and  therefore  the  percentage,  of  the  corresponding  substance  in 
the  sample  can  be  calculated. 

The  substance  to  be  weighed  is  in  most  cases  separated  from 
solution  by  precipitation,  though  in  many  instances  it  is  de- 
posited upon  a  weighed  cathode  or  anode  by  electrolysis.  Some- 
times it  is  separated  from  other  substances  by  extraction  with  a 
solvent,  and  sometimes  in  the  form  of  a  gas,  the  weight  of  the 


2  QUANTITATIVE  CHEMICAL  ANALYSIS 

gas  being  determined  either  by  absorbing  it  in  a  weighed  quan- 
tity of  some  substance  and  noting  the  increase,  or  by  noting  the 
decrease  in  weight  due  to  the  removal  of  the  gas  alone. 

In  a  volumetric  analysis  a  weighed  sample  is  also  taken,  but 
the  quantity  of  the  substance  to  be  determined  is  arrived  at 
by  causing  a  definite  reaction  to  take  place,  the  reagent  being 
added  from  a  burette,  as  a  solution  of  known  concentration. 
This  operation  is  called  titration.  From  the  volume  of  the  solu- 
tion added,  it  is  easy  to  calculate  the  weight  of  the  substance 
present  in  the  sample. 

In  many  instances,  it  is  necessary  in  volumetric  analysis  also 
to  separate  the  substance  to  be  determined  from  interfering 
substances  present  with  it  in  the  sample ;  but,  instead  of  finally 
weighing  it,  the  substance  is  again  brought  into  solution,  in 
suitable  form,  and  its  quantity  estimated  by  titration. 

In  order  to  illustrate  the  two  methods,  let  us  consider  the 
determination  of  silver  in  a  silver  coin. 

(a)  Gravimetric  Method.  The  weighed  sample  is  dissolved 
in  nitric  acid,  the  solution  diluted,  and  the  silver  separated  from 
copper,  by  precipitation  as  insoluble  silver  chloride,  with  dilute 
hydrochloric  acid.  The  precipitate  is  filtered  off,  washed,  dried, 
and  weighed.  From  its  weight,  the  weight  of  silver  is  calcu- 
lated, as  follows : 

Ag 

7— p-  Xwt.  of  precipitate =wt.  of  silver. 

A    J     r  wt.  of  silver  ^^  ^    r    m 

And,  of  course,  -— — -z r-  Xioo=per  cent  of  silver. 

wt.  of  sample 

(b)  Volumetric  Method.  The  weighed  sample  is  dissolved 
in  nitric  acid,  diluted  as  before,  and  the  silver  converted  into 
the  insoluble  chloride  by  the  gradual  addition,  from  a  burette, 
of  a  solution  of  sodium  chloride  of  known  concentration.  As 
soon  as,  after  stirring  each  time  and  allowing  the  precipitate  to 
settle,  the  first  drop  is  added  which  fails  to  induce  further  pre- 
cipitation, the  reaction  is  known  to  be  complete ;  and  the  number 
of  cubic  centimeters  required,  multiphed  by  the  silver  equivalent 


INTRODUCTION  3 

of  the  sodium  chloride  solution  per  cubic  centimeter,  gives  di- 
rectly the  weight  of  silver  in  the  sample. 

B.   GENERAL  REMARKS  CONCERNING 
QUANTITATIVE  WORK 

Physical  Conditions.  The  inside  and  top  of  the  desk,  the  re- 
agent bottles,  and  all  apparatus  should  be  kept  clean  and  ready 
for  use.  Vessels  for  use  in  analytical  work  should  be  perfectly 
clean,  and  the  outside  of  glassware  should  be  wiped  dry  before 
use,  especially  if  it  is  to  be  heated  over  a  burner,  or  on  the  hot 
plate.  Flasks,  funnels,  etc.  may  be  freed  from  inside  grease 
films  by  rinsing,  or  by  longer  contact,  with  either  a  solution  of 
sodium  hydroxide  in  alcohol,  or  with  sulphuric  acid-dichromate 
cleaning  solution.  (For  the  preparation  of  the  latter,  see  the 
Appendix.) 

Reagents  and  Glassware.  One  of  the  greatest  hindrances 
to  exact  work  hes  in  the  difficulty  of  securing  satisfactory  re- 
agents. And,  in  addition  to  this,  much  of  the  glassware  on  the 
market  is  of  an  inferior  grade,  and  not  at  all  suited  for  analyti- 
cal work. 

The  habit  of  carefully  testing  reagents,  including  distilled 
water,  cannot  be  too  early  acquired ;  constant  vigilance  should 
be  exercised  in  guarding  against  the  presence  of  impurities  which 
would  vitiate  the  work  under  way.  As  is  generally  known,  a 
"  C.  P."  label  is  no  guaranty  whatever  of  the  purity  of  a  re- 
agent, and  the  "  guaranteed  "  or  "  analyzed  "  reagents,  sold  at 
higher  prices,  are  at  times  inferior  to  products  for  which  no 
special  claim  is  made. 

Very  pure  acids  are  obtainable  in  the  market,  and,  in  general, 
these  need  no  redistillation.  But,  owing  to  its  basic  nature, 
ammonia  ought  to  be  redistilled  at  short  intervals,  after  a  pre- 
liminary treatment  with  slacked  lime  to  remove  absorbed 
carbonic  acid ;  to  prevent  contamination  of  the  solution  through 
its  solvent  action  upon  glass,  stock  bottles  for  ammonia  may  be 
coated  inside  with  ceresin. 


4  QUANTITATIVE  CHEMICAL  ANALYSIS 

Owing  to  the  solvent  action  on  glass  of  many  solutions  of 
solid  reagents,  these  should  be  prepared  frequently  and  in  small 
quantities ;  or,  better,  the  solid  should  be  dissolved  as  wanted. 
This  applies  particularly  to  such  reagents  as  ammonium  oxalate, 
microcosmic  salt  and  sodium  phosphate,  alkaline  magnesia 
mixture,  ammonium  carbonate,  etc. 

The  stopper  of  a  reagent  bottle  should  always  be  held  in  the 
fingers  until  returned  to  the  bottle.  This  will  prevent  con- 
tamination, whether  due  to  an  interchange  of  stoppers,  or  to  some 
other  cause.  The  contents  of  such  bottles  must  of  course  be 
protected  at  all  times  from  contamination. 

Utilization  of  Time.  Economy  of  time  in  the  laboratory  is 
best  insured  through  a  preliminary  study  of  the  work  in  pros- 
pect, followed  by  a  plan  for  its  prompt  execution.  It  may  be 
stated  in  general  that  several  analyses  should  always  be  under 
way,  and  that  no  vessel  should  ever  be  left  unlabelled.^  In  this 
way,  the  analyst  will  always  have  something  to  occupy  his  time. 
At  the  outset,  of  course,  the  student  cannot  be  expected  to  ap- 
proach this  degree  of  efficiency;  nevertheless,  even  here,  time 
can  'be  either  wasted  or  utiHzed,  —  e.g.  the  student  should 
bear  in  mind  that  the  time  required  for  the  solution  of  several 
samples,  or  for  the  filtration  or  evaporation  of  several  solu- 
tions, is  the  same  as  that  required  in  the  case  of  a  single  one. 

Accuracy.  Quantitative  work,  to  be  successful,  must  be  per- 
formed without  the  slightest  loss  of  material,  as  well  as  without 
gain  of  material  from  the  outside.  Whenever  solutions  are  to 
be  boiled,  therefore,  or  in  dissolving  substances  for  analysis, 
especially  if  finely  divided,  or  if  effervescence  is  at  all  likely, 
the  vessels  should  be  kept  covered ;  also,  in  the  evaporation  of 
liquids,  even  on  the  steam  bath,  large  watch  glasses  should  be 
supported  above  the  vessels,  in  order  to  prevent  the  entrance  of 
dust,  while  at  the  same  time  allowing  the  vapor  to  escape  freely. 

*  In  certain  cases,  it  is  customary  to  keep  the  record  in  terms  of  serial  numbers. 
Platinum  crucibles,  for  example,  sliould  have  stamped  upon  them  serial  numbers, 
and  under  these  numbers  records  should  be  kept  in  the  notebook  of  their  contents. 


INTRODUCTION  5 

It  is  of  course  indispensable  that  the  analyst  should  be  pos- 
sessed of  a  certain  degree  of  dexterity  in  the  performance  of  the 
quantitative  operations.  But,  while  a  few  persons  seem  to  have 
been  born  with  this  skill,  the  majority  can  acquire  it  only  through 
careful  and  persistent  application.  The  student  who  finds  him- 
self at  first  unable  to  do  good  work,  therefore,  should  not  try 
to  hide  the  fact,  but  should  patiently  endeavor  to  overcome  the 
handicap ;  nothing  can  he  more  fatal  to  successful  scientific  work 
than  a  lack  of  personal  integrity. 

All  determinations  should  be  made  in  duplicate,  and  in  general 
the  results  should  closely  agree.  "  Check  results,''  however, 
do  not  necessarily  indicate  accuracy ;  they  depend  very  largely 
upon  identical  conditions  in  the  two  analyses,  and  are  possible 
even  in  the  case  of  inaccurate  methods.  A  failure  to  obtain 
check  results,  however,  is  conclusive  in  its  testimony ;  one  re- 
sult is  certainly  wrong,  and  the  other  is  little  worthy  of  con- 
fidence. 

It  is  often  thought  that  no  operation  in  an  analysis  need  be 
carried  out  with  more  exactitude  than  the  one  which  necessarily 
involves  the  greatest  error ;  that  if  a  given  step  cannot  be  exe- 
cuted without  an  uncertainty  of  o.i%,  it  would  be  useless  to 
avoid  errors  of  that  magnitude  in  other  stages  of  the  process. 
If  carried  to  its  logical  conclusion,  this  would  imply  that  a 
method  capable  of  furnishing  results  within,  say,  0.15%  of  the 
true  value,  might  just  as  well  be  allowed  to  yield  results  a  few 
tenths  of  a  per  cent  farther  off.  Of  course,  a  compensation  of 
errors  might  take  place,  but  this  is  not  always  even  probable. 
Nevertheless,  if  an  error  of,  say  0.15%  cannot  be  avoided,  it 
would  be  inadvisable  to  devote  much  labor  towards  the  reduc- 
tion of  a  very  much  smaller  individual  error  also  involved  in  the 
process. 

Laboratory  Records.  Notebooks  should  contain  the  analyti- 
cal data,  systematically  entered  at  the  time  of  observation, 
and,  in  addition,  an  account  of  any  unexpected  development 
in  the  analysis.    Since  a  neat  and  intelHgible  record  is  of  the 


6  QUANTITATIVE  CHEMICAL  ANALYSIS 

greatest  importance,  the  following  sample  page  is  offered  in  the 
way  of  suggestion.  In  the  calculation  of  the  result  from  the 
analytical  data,  it  is  useless  to  proceed  more  than  one  place 
beyond  the  last  that  may  reasonably  be  considered  correct; 
according  to  the  data  in  this  case,  it  is  not  certain  whether  the 
first  figure  after  the  deeimal  should  be  4,  3,  or  2,  —  nothing, 
therefore  can  be  known  concerning  the  second  figure,  and  it  is 
not  included  in  the  report  of  the  mean. 

Determination  of  Chlorine  in  a  Soluble  Chloride  —  Sept.  21-28 


I 

n 

Sample  Tube,  etc. 

«.4237 

8.2377 

Tube  minus  Sample 

8.2377 

8.0198 

Wt.  of  Sample 

0.1860 

0.2179 

Wt.  of  Crucible  No. 

(7)  5.3588 

(8)  5.0072 

Crucible-hAgCl,  ist  time 

S.7830 

5.5024 

2d  time 

S.7828 

5.5023 

Wt.  of  Crucible 

5.3588 

5.0072 

Wt.  of  AgCl 

0.4240 

0.4951 

Per  cent  of  Chlorine 

56.40 

56.21 

Mean  Value = 

=  56.3% 

C.  THE  OPERATIONS  OF  ANALYTICAL  CHEMISTRY 

The  chief  operations  involved  in  analytical  work  which  can 
be  profitably  discussed  at  this  point  are  weighing,  precipitation, 
filtration,  and  the  washing  of  precipitates,  the  drying  and  ig;tii- 
tion  of  precipitates,  the  evaporation  of  liquids,  and  the  volumet- 
ric measurement  of  liquids. 

These  operations  will  be  described  in  the  following  sections, 
which  should  be  studied  carefully  by  the  beginner.  The  student 
should  from  the  start  try  to  appreciate  the  importance  of  pay- 
ing great  attention  to  details  and  of  closely  observing  the  condi- 
tions necessary  for  the  greatest  accuracy. 


INTRODUCTION  7 

I.  WEIGHING! 

Standards  of  Mass.  The  fundamental  standard  of  mass 
adopted  by  the  United  States  is  a  cylinder  of  platinum-iridium 
kept  at  the  International  Bureau  of  Weights  and  Measures 
near  Paris.  Two  authentic  copies  of  this  standard,  of  the  same 
form  and  composition,  are  kept  in  a  vault  at  the  National 
Bureau  of  Standards,  and  they  are  used  only  when  needed  to 
verify  the  secondary  standards  of  the  Bureau. 

Mass  standards  are  usually  called  "  weights,"  but  it  should 
be  realized  that,  while  the  weight  of  a  body  depends  upon  its 
attraction  by  the  earth,  the  mass  of  a  body  is  a  property  in- 
herent in  the  body  itself.  Since,  however,  the  masses  of  bodies 
are  proportional  to  their  weights  at  the  same  locality,  we  may 
compare  the  masses  by  making  a  comparison  of  the  respective 
forces  of  gravity  on  the  bodies  under  consideration.  The  pur- 
pose of  weighing,  then,  is  to  compare  the  quantity  of  matter  in  a 
specific  object  with  the  quantity  of  matter  in  a  given  standard 
—  a  gram  or  kilogram  weight.  The  comparison  is  made  on  the 
balance,  by  means  of  the  lever  principle,  by  suspending  the  ob- 
ject at  one  end  of  a  beam,  and  the  weights  at  the  other  end  of 
the  beam,  the  beam  being  virtually  a  kind  of  lever. 

The  Balance.  The  beam  of  the  balance  is  supported  on  a 
central  knife-edge,  usually  of  agate,  which  rests  upon  a  plane 
agate  plate;  and  two  pans  for  supporting  the  masses  to  be 
compared  are  vertically  suspended  from  stirrups,  each  of  which 
has  an  agate  bearing  which  rests  on  a  knife-edge  fixed  at  one 
extremity  of  the  beam.  The  arms  of  the  balance  are  so  gradu- 
ated that  a  rider  (of  known  weight)  can  be  placed  on  the  beam  at 
any  required  distance  from  the  central  knife-edge. 

If  the  three  knife-edges  are  allowed  to  press  continually  upon 
their  agate  bearings,  they  soon  become  blunted,  and  wear  fur- 
rows in  the  bearings.     In  order  to  prolong  the  life  of  the  knife- 

^  For  more  detailed  information  on  the  subject  of  weights  and  weighing,  the 
student  is  referred  to  Circular  of  the  Bureau  of  Standards,  No.  3  (3d  edition,  1918), 
from  which  much  of  the  following  is  taken. 


8  QUANTITATIVE  CHEMICAL  ANALYSIS 

edges  and  bearings,  the  balance  is  provided  with  a  "release" 
which  separates  the  knife-edges  from  their  bearings  when  the 
balance  is  not  in  use.  If  the  balance  shows  signs  of  stiffness 
in  the  motion  of  beam  and  pans,  the  fault  should  be  investigated 
at  once.  The  defect  may  be  due  to  an  accumulation  of  dust 
between  the  knife-edges  and  their  bearings ;  to  the  blunting  of 
the  knife-edges;  or  to  the  wearing  of  furrows  in  the  bearings. 
To  prevent  the  accumulation  of  dust,  and  also  to  prevent  the 
interference  of  air  currents  while  weighing,  the  balance  is  in- 
closed in  a  glass  case. 

In  order  to  render  small  movements  of  the  beam  perceptible, 
there  extends  downwards  from  its  center  a  long  pointer  which 
multiplies  the  rotational  displacement.  When  equilibrium  is 
established,  the  lower  end  of  the  pointer  should  come  to  rest  in 
front  of  the  zero  of  a  scale  which  is  located  immediately  behind 
this  end. 

The  conditions  which  must  be  satisfied  by  a  good  balance  are : 
(i)  The  balance  must  be  consistent.  It  must  give  the  same 
result  in  successive  weighings  of  the  same  body.  This  condition 
depends  upon  the  trueness  of  the  knife-edges.  (2)  The  balance 
must  be  accurate.  At  rest  the  beam  must  be  horizontal  when 
the  pans  are  empty,  and  when  equal  weights  are  placed  upon 
the  pans.  This  condition  depends  upon  the  equality  of  the 
two  arms.  (3)  The  balance  must  be  stable.  The  beam  after 
being  displaced  from  its  horizontal  position  must  return  to  its 
horizontal  position.  This  condition  depends  upon  the  adjust- 
ment of  the  center  of  gravity.  (4)  The  balance  must  be  sensi- 
tive. It  must  show  even  a  very  small  inequahty  in  the  two 
masses  on  the  scale  pans.  This  condition  depends  largely  upon 
the  length  of  the  arms.  (5)  The  balance  must  oscillate  with 
reasonable  rapidity.  Short  beams  oscillate  more  rapidly  than 
long  ones. 

The  analytical  balance  will  perform  excellent  service  under 
the  proper  conditions,  but  great  care  in  its  use  is  essential  if  its 
accuracy  is  to  be  relied  upon.     It  should  be  located  in  a  room 


INTRODUCTION  9 

that  is  free  from  dust  and  fumes,  and  should  stand  upon  a  sup- 
port that  is  free  from  shocks  and  vibrations. 

The  Use  and  Care  of  the  Analytical  Balance.  The  following 
rules  embody  the  main  points  to  be  observed  in  the  use  and 
care  of  a  balance. 

(i)  Each  student  must  feel  a  personal  responsibility  for  the 
proper  use  of  his  balance;  the  carelessness  of  any  one  may 
render  inaccurate  the  work  of  all  who  use  the  same  balance. 

(2)  Before  use,  with  clean  pans,  the  adjustment  of  the  bal- 
ance should  be  tested  by  the  student. 

The  balance  is  properly  adjusted  only  if  the  following  con- 
ditions are  fulfilled:  {a)  The  spirit  level  or  plumb  bob  inside 
the  balance  case  should  show  that  the  balance  is  level ;  (6)  the 
mechanism  for  raising  and  lowering  the  beam  should  work 
smoothly ;  {c)  the  pan  arrests  should  just  touch  the  pans  when 
the  beam  is  lowered ;  {d)  the  pointer  should  rest  at  zero  when 
the  beam  is  raised ;  and  also  when  it  is  lowered,  with  the  pans 
supported;  and  (e)  the  pointer  should  swing  equal  distances 
on  either  side  of  the  zero  when  the  beam  is  set  in  motion  with- 
out any  load  on  the  pans.  In  the  latter  case,  if  the  variation 
does  not  exceed  two  divisions  on  the  scale,  it  is  sufficient  to 
allow  for  the  small  zero  error,  without  adjustment. 

(3)  The  beginner  should  not  attempt  to  make  adjustments 
himself,  but  should  apply  to  the  instructor  in  charge. 

(4)  In  order  to  set  the  beam  in  motion,  it  should  first  be 
lowered  so  that  the  pans  rest  upon  the  pan  arrests ;  the  latter 
are  then  cautiously  lowered,  and,  in  case  the  beam  fails  to 
swing,  it  can  be  set  in  motion  by  means  of  the  rider.  There  is, 
however,  a  "  trick  "  in  lowering  the  pan  supports  so  that  the 
oscillations  of  the  pointer  will  have  the  required  amplitude.^ 

The  pans  should  be  arrested  and  the  beam  raised  before  any 
change  is  made  in  the  load  or  weights  on  the  pans  except  in  the 

1  It  is  not  necessary  to  employ  very  long  swings ;  an  amplitude  of  two  scale 
divisions  is  probably  sufficient,  even  in  very  exact  work.  See  in  this  connection 
Horace  L.  Wells,  Jour.  Amer.  Chem.  Soc,  vol.  42,  p.  411  (1920). 


10  QUANTITATIVE  CHEMICAL  ANALYSIS 

case  of  the  small  fractional  weights,  when  it  is  only  necessary 
to  arrest  the  pans.  The  object  to  be  weighed  and  the  heavy 
weights  should  be  placed  in  the  middle  of  their  respective  pans, 
since  a  heavy  load  near  the  edge  of  a  pan  is  apt  to  cause  trouble- 
some oscillations. 

tV  The  beam  and  stirrups  should  never  be  left  upon  their  knife- 
edges,  and  the  motion  of  the  beam  should  be  arrested  only  by 
means  of  the  pan  arrests,  and  only  when  the  pointer  is  passing 
the  center  of  the  scale;  otherwise  the  knife-edges  and  their 
agate  bearings  are  subjected  to  an  unnecessary  strain. 

(5)  The  weights  should  be  cared  for  as  well  as  the  balance. 
They  should  be  handled  carefully,  and  only  with  the  forceps 
provided;  they  should  never  be  touched  with  the  fingers.  In 
weighing,  the  weights  should  always  be  placed  upon  the  same 
pan,  and  they  should  be  taken  in  the  order  in  which  they  occur 
in  the  box,  the  larger  ones  first ;  and  the  weight  of  the  object 
should  be  recorded  by  noting  the  vacant  spaces  in  the  box.  The 
record  so  obtained  should  be  checked  as  the  weights  are  re- 
moved from  the  pan.  In  this  way  errors  are  not  likely  to 
occur. 

(6)  Analytical  samples  should  not  be  placed  directly  upon 
the  balance  pan,  and  the  object  to  be  weighed  should  not  be 
warmer  or  colder  than  the  air  in  the  balance  case.  Currents 
of  hot  air  will  tend  to  buoy  up  one  arm  of  the  balance,  and  also 
to  cause  that  arm  to  expand  in  length.^  If  the  object  is  colder 
than  the  atmosphere  of  the  balance  case,  moisture  may  con- 
dense on  its  surface.  If  the  body  to  be  weighed  is  likely  to  be 
electrified  {e.g.  a  glass  weighing  tube),  it  should  be  allowed  to 
stand  for  some  time  after  it  has  been  wiped,  before  weighing. 

(7)  The  balance  case  should  be  closed  while  weighing  with 
the  rider,  so  as  to  avoid  currents  of  air. 

As  soon  as  the  object  is  apparently  balanced  by  the  weights 
the  beam  should  be  raised  and  again  lowered  into  place,  and  the 

*  For  instance,  a  platinum  crucible  which  appeared  to  weigh  20.649  g.  when 
warm,  weighed  20.6920  g.  when  cold. 


INTRODUCTION  il 

observation  repeated.    This  insures  the  proper  alignment  of 
the  beam  and  pans. 

(8)  In  using  weighing  bottles  or  tubes,  the  vessel  should  be 
weighed  together  with  its  contents.  The  sample  should  then  be 
removed  without  loss,  and  the  vessel  and  contents  again  weighed. 
The  difference  in  weight  indicates  the  quantity  taken.  The 
weight  of  a  tube,  recorded  at  some  previous  time,  should  always 
be  confirmed  before  weighing  out  a  new  sample  from  it. 

(9)  Errors  in  weighing  should  fall  well  within  the  Umits  of 
the  experimental  error  due  to  the  analytical  operations.  If, 
for  example,  an  error  of  o.ooi  g.  were  made  in  weighing  out  a 
gram  sample  of  fireclay  containing  0.25%  of  MgO,  the  result- 
ing error  in  the  determination  of  the  magnesia  could  be  no 
greater  than  0.1%  of  its  value,  which  is  negligibly  small.  A 
I  mg.  error,  however,  made  in  weighing  the  0.0069  g.  of  Mg2P207 
would  involve  an  error  of  over  14%  in  the  magnesia  value,  and 
this  would  be  intolerable. 

(10)  Finally,  if  anything  at  all  appears  to  be  the  matter  with 
a  balance,  the  instructor's  attention  should  at  once  be  called 
to  the  fact. 

Determination  of  the  Rest-point.  The  beam  and  stirrups 
are  first  lowered  upon  their  knife-edges  by  slowly  turning  to  the 
left  the  milled  head  at  the  front  of  the  balance  case ;  then  the 
pans  are  released  by  gently  pressing  inwards  the  small  button, 
also  at  the  front  of  the  case;  and,  with  the  beam  swinging 
smoothly,  a  consecutive  record  is  made  of  the  number  of  scale 
divisions  traversed  by  the  pointer  on  either  side  of  the  center. 
The  swings  to  the  left  are  recorded  as  negative  numbers  and 
those  to  the  right  as  positive  numbers ;  in  the  determination  of 
the  rest-point,  one  more  reading  must  be  made  on  one  side  than 
on  the  other,  and  all  of  the  readings  must  be  consecutive.  Upon 
dividing  by  2  the  algebraic  sum  of  the  averages  for  the  two 
sides,  the  quotient  is  the  rest-point  of  the  balance  for  the  case 
under  consideration,  i.e.  the  position  on  the  scale  at  which  the 
pointer  would  finally  come  to  rest. 


12 


QUANTITATIVE  CHEMICAL  ANALYSIS 


Example: 

Left 

Right 

-4.8 
-4.6 
-44 

+2.7 
+2.5 

Average:                 —4.6 

Average:                +2.6 

Rest-point  =  — 1,0. 

Two  methods  of  procedure  are  now  open  to  the  operator. 
He  may  either  make  his  weighings  with  reference  to  this  ob- 
served rest-point  or  he  may  adjust  the  balance  so  that  the  ob- 
served rest-point  is  the  actual  zero  of  the  scale.  The  first 
method  is  preferable,  unless  the  rest-point  is  more  in  error  than 
one  scale  division.  The  rest-point  is  apt  to  change,  and  it  must 
be  determined  each  day,  or  even  more  often. 

Methods  of  Weighing.  Weighings  smaller  than  0.005  g-  (or 
o.oi  g.)  are  made  with  the  rider.  When  the  arms  are  divided  into 
five  divisions,  a  5 -milligram  rider  is  used ;  in  general,  the  rider 
should  weigh  as  many  milligrams  as  there  are  large  divisions 
between  the  central  knife-edge  and  the  right-hand  stirrup  sup- 
port. Each  large  division  on  the  beam  then  corresponds  to  a 
milligram. 

Ordinary  Method,  The  object  to  be  weighed  is  placed  upon 
the  left-hand  pan  of  the  balance  and  weights  upon  the  right- 
hand  pan,  until,  finally,  the  further  addition  of  5  mg.  (or  10  mg.) 
more  than  counterbalances  the  object.  This  weight  is  then  re- 
moved, the  balance  case  closed,  and  the  rider  adjusted  so  that  the 
pointer  swings  equal  distances  on  either  side  of  the  rest-point. 
This  method  of  weighing  is  very  common,  and  it  is  sufficiently 
accurate  for  ordinary  analytical  work.  If  necessary,  the  rest- 
point  of  the  unloaded  balance  should  be  determined  before  each 
weighing. 

In  special  cases,  as  in  the  calibration  of  a  set  of  weights,  it  is 
important  to  make  more  accurate  weighings.    It  is  here  best  to 


INTRODUCTION  13 

employ  the  method  of  weighing  by  the  use  of  deflections.  Though 
this  method  may  appear  to  be  laborious,  the  labor  is  more  ap- 
parent than  real.  The  sensitiveness  of  the  balance  may  be 
found  by  adding  a  small  weight  to  one  of  the  pans  and  noting 
how  far  the  rest-point  is  deflected  from  its  former  position. 
Then,  instead  of  adjusting  small  weights  until  the  rest-point  is 
brought  to  the  proper  place,  we  merely  note  the  deflection 
from  this  point  and  calculate  from  the  sensitiveness  the  weight 
that  would  be  needed  to  bring  the  rest-point  to  the  desired  posi- 
tion. The  sensitiveness  will,  in  general,  be  different  at  differ- 
ent loads,  and,  especially  with  a  very  sensitive  balance,  it  must 
be  redetermined  from  time  to  time.  For  very  accurate  work, 
it  is  advisable  to  determine  the  sensitiveness  at  each  weigh- 
ing. 

Method  of  Weighing  by  the  Use  of  Deflections,  (a)  The  rest- 
position  of  the  unloaded  balance  is  determined,  as  already 
described.     Let  us  suppose  this  to  be  at  -f-o.i. 

(b)  The  deflection  of  the  rest-point  per  milligram,  or  the 
sensitiveness  of  the  loaded  balance,  is  determined.  The  object 
to  be  weighed  is  placed  upon  the  left  pan,  the  weights  on  the 
right  pan.  When  the  weights  have  been  adjusted  so  that  an 
additional  0.005  g.  (or  o.oi  g.)  would  be  too  much  (e.g.  weights 
=  11.216  g.),  the  balance  case  is  closed,  and  the  rider  adjusted 
until  the  pointer  swings  on  both  sides  of  the  zero  of  the  scale. 
The  rest-point  is  then  found ;  for  example,  at  H-o.8.  The  rider 
is  moved  one  milligram  division,  in  this  case  to  the  right,  and  the 
rest-point  again  determined;  at,  say,  —2.1.  That  is,  the 
rest-point  is  deflected  -fo.8  — (  — 2.i)  =  2.9  divisions  by  i  milli- 
gram, in  the  case  of  that  particular  load. 

(c)  The  weight  of  the  object  is  calculated.  In  this  case, 
the  object  weighs  ii.2i6H-:x;  g.  The  rest-point  of  this  load 
is  displaced  0.8—0.1  =  0.7  scale  division.  Since  2.9  scale  divi- 
sions correspond  to  i  milligram,  0.7  scale  division  will  correspond 

to-^=o.24  mg.     Hence  the  weight  of  the  body  is  ii.2i6-|- 
2.9 


14  QUANTITATIVE  CHEMICAL  ANALYSIS 

0.00024=11.21624  g.  These  calculations  may  be  summarized 
in  the  formula  _ 

Correction  =  + — r-^g-> 
a  —  o 

in  which  z  represents  the  rest-point  of  the  unloaded  balance, 
a,  the  rest-point  with  not  quite  enough  weight  on  the  right  pan ; 
and  hj  the  rest-point  with  a  milligram  more  on  the  right  pan 
than  corresponds  to  a. 

Analytical  balances  will  rarely  indicate  with  certainty  less 
than  o.oooi  g.  Hence,  although  the  weight  may  be  calculated 
as  above  to  the  fifth  decimal,  it  should  generally  be  rounded  off 
by  dropping  the  fifth  decimal  and  raising  the  fourth  decimal 
one  unit  when  the  dropped  figure  exceeds  5. 

In  certain  cases,  as  in  the  calibration  of  volumetric  measuring 
apparatus,  it  is  necessary  for  the  weight  found  to  be  independent 
of  any  inequality  in  length  in  the  beam  arms.  In  such  cases, 
and  in  the  determination  of  absolute  weights  (reduction  to 
weights  in  vacuo),  one  of  the  following  methods  should  be  used. 

Method  of  Weighing  by  Transposition.  When  the  arms  of  a 
balance  are  nearly  equal,  the  method  of  transposition  furnishes 
a  more  accurate  comparison  than  that  of  substitution  (see  be- 
low). It  requires  about  the  same  amount  of  time  for  the  obser- 
vations, but  the  true  weight  is  not  shown  in  so  direct  a  manner. 
Before  and  after  the  transposition,  the  rest-point  of  the  loaded 
balance  is  noted  and  the  added  weight  that  would  be  needed 
to  bring  the  rest-point  to  the  desired  position  is  determined  as 
already  described. 

Having  weighed  the  object  first  in  one  pan,  then  in  the  other, 
if  we  let  W  be  the  true  weight,  a  the  weights  required  to  counter- 
balance the  object  when  it  is  on  the  left  pan,  and  b  the  weights 
required  when  the  object  is  on  the  right  pan,  then,  according  to 
the  principle  of  moments : 

Wl   =  ar;  andW    =Wr, 

That  is,  Wlr=  ablr,  or  W^  =ab: 

whence  W=  "y/ab. 


INTRODUCTION  15 

Therefore  the  true  weight  is  the  square  root  of  the  product  of 
the  two  observed  weights. 

Weighing  by  transposition  is  recommended  for  work  of  high 
precision  in  which  it  is  also  desirable  to  calculate  the  rest-point 
from  several  swings  of  the  pointer.  In  other  cases,  substitution 
is  generally  to  be  preferred. 

Method  of  Weighing  by  Substitution.  Here  the  object,  placed 
on  the  right-hand  pan,  is  approximately  balanced  by  a  suitable 
tare  (weights,  wire,  beaker  containing  shot,  etc.)  on  the  left- 
hand  pan,  and  the  rest-point  determined.  The  object  is  then  re- 
moved, and  weights  are  added  in  its  place  until  the  rest-point 
is  restored  to  its  former  position.  These  weights  are  necessarily 
the  same  in  value  as  the  object  for  which  they  substitute,  irre- 
spective of  differences  in  the  arms. 

The  Calibration  of  a  Set  of  Weights.^  Fairly  accurate  weights 
can  readily  be  purchased,  and  for  most  analytical  work  the  in- 
accuracies of  the  better  class  of  weights  are  negligibly  small  in 
comparison  with  the  errors  of  experiment,  and  the  imperfections 
in  the  analytical  processes. 

An  analyst,  however,  should  know  that  his  weights  are  suffi- 
ciently accurate,  and  for  this  reason  he  should  calibrate  the 
weights.  The  errors  in  weighing  due  to  imperfections  in  the 
weights  can  easily  be  reduced  to  o.oooi  g.  The  weights  should 
be  tested  at  periodic  intervals,  say  once  or  twice  a  year,  depend- 
ing upon  the  frequency  with  which  they  are  used. 

In  special  cases,  e.g.  in  the  calibration  of  volumetric  apparatus, 
absolute  weights  may  be  required,  but  for  general  analytical 
work  these  are  not  necessary.  If  the  weights  are  consistent 
with  one  another,  their  absolute  values  have  no  influence  upon 
the  accuracy  of  an  analysis. 

Duplicate  and  triplicate  weights  of  a  set  should  be  so  marked 
that  they  can  be  readily  distinguished.  A  satisfactory  method 
is  to  use  one  and  two  conspicuous  dots  on  duplicates,  and  one, 

*  In  this  connection,  see  Circular  of  the  Bureau  of  Standards,  No.  j,  and  also 
T.  W.  Richards,  Jour.  Amer.  Chem.  Soc,  vol.  22,  p.  144  (1900). 


i6  QUANTITATIVE  CHEMICAL  ANALYSIS 

two,  and  three  such  dots  on  triplicates,  —  stamped  by  means 
of  a  small  punch.  The  designations  of  the  weights  being  tested 
are  conveniently  inclosed  in  parentheses,  to  prevent  their  con- 
fusion with  the  numerical  data  for  use  in  the  calculations.  And 
it  is  customary  to  express  the  results  in  terms  of  corrections  to 
be  applied  in  the  use  of  the  weights,  rather  than  in  terms  of  the 
total  values;  the  plus  sign  for  a  correction  indicating  that  the 
mass  of  a  weight  is  greater  than  its  nominal  value,  the  minus 
sign  that  it  is  less.  In  using  the  weights,  these  signs  are  then 
treated  in  the  ordinary  algebraic  manner.  For  example,  if 
(100  =  9.9987,  (10")  =  lO-oooS,  and  (5)  =  5.0002,  the  respective 
corrections  are  —1.3,  +0.8,  and  +0.2  mg. ;  the  nominal  value 
of  the  weights  is  25.0000  g.,  but  their  corrected  value  is  25.0000+ 
(-1.3+0.8+0.2)  mg.,  or  24.9997  g. 

The  weights  are  best  calibrated  according  to  their  apparent 
masses  as  determined  by  comparison  with  brass  standards  in 
air;  that  is,  in  terms  of  "weight  in  air  against  brass."  This 
method  is  advisable,  owing  to  the  prevailing  use  of  small  plati- 
num and  aluminum  weights  in  connection  with  the  larger  weights 
of  brass. 

In  the  calibration  of  analytical  weights,  the  rest-point  should 
always  be  calculated  from  several  swings  of  the  pointer,  and  the 
individual  calculations  should  be  extended  to  five  decimals. 
If  the  beam  arms  do  not  differ  in  length  by  more  than  0.001%, 
the  simple  method  of  deflections  may  be  used ;  otherwise  it  is 
necessary  to  make  use  of  transposition  or  substitution.  The 
substitution  method  involves  less  work  in  calculating  than  that 
of  transposition,  but  if  it  is  used  it  is  desirable  to  have  a  second 
set  of  weights  to  furnish  the  tares. 

In  the  case  of  an  ordinary  analytical  set  of  weights,  from  20  g. 
to  5  mg.,  the  following  comparisons  should  be  made,  with  the 
use  of  the  rider,  as  in  ordinary  weighings. 


INTRODUCTION 


17 


The  Small  Weights^ 


The  Larger  Weights 


(500)   against  2(500) 

(20)   against    (ioO+(io") 

(200)   against    (iooO+(ioo") 

(loO  against    (lo'O 

(looO  against    (100") 

(loO  against  5(io) 

(100')  against  2(ioo) 

(S)     against  5(5) 

(50)     against  2  (50) 

(2)     against    (i')+(i") 

(20)     against    (ioO+(io") 

(i')    against    (i") 

(10')    against    (lo'O 

(i')    against    (i'") 

(10')    against  2  (10) 

(s)       against  (Rider  at  5) 

One-gram  piece,  (i'),  against3(i) 


D 

Ten-gram  piece,  (10'),  against  an 
absolute  (known)  standard  of  the 
same  nominal  value. 


The  weight  of  each  of  the  smaller  pieces  is  calculated  from  the 
data  in  terms  of  one  of  them,  say  the  ten-milligram  piece  (10'), 
as  a  standard ;  and,  in  the  same  way,  that  of  each  larger  piece 
in  terms  of  one  of  them,  as  the  ten-gram  piece  (10').  Then, 
in  order  to  express  all  the  weights  of  the  set  in  terms  of  a  com- 
mon standard,  the  values  of  the  fractional  weights  (in  terms  of 
their  standard)  are  added,  and  the  result  compared  with  their 
total  weight  as  found  in  terms  of  a  one-gram  piece,  say  (i')- 

Suppose,  for  example,  the  values  of  the  small  weights  are: 
(5oo)  =  499.03,  (200)  =  200.07,  (100)'=  100.10,  (loo'O  =  100.00, 
(50)  =  49-87,  (20)  =  19.98,  (ioO  =  10.00,  (io")=  10.03,  (5)  = 
5.00,  and  (Rider  at  5)  =  5.00 — i.e.  2 (i)  =  0.99908  ;  while,  in  terms 
of  the  one-gram  piece  (i'),  their  collective  weight  is  i. 00165. 
Then,  in  order  to  express  the  values  in  terms  of  (i'),  each  of  them 

should  be  multiplied   by  — ^=i.oo2';7.      If,  however,    in 

^  -^    0.99908  ^/  ' 

terms  of  the  ten-gram  piece  (10'),   (10  =  0.99954,  the  provisional 

*  The  designations  S  (500),  etc.,  are  here  used  to  represent  collective  weights  of 
the  nominal  values  indicated;  in  this  instance,  (200)-}- (100')  -h(ioo")+(so)-{-  (20) 
-h(io')+(io")+(s)  +  (Rider  at  5). 


i8  QUANTITATIVE  CHEMICAL  ANALYSIS 

value  of  each  fractional  piece  may  be  multiplied  directly  by 

1.00165X0.99954  ^  ^,    .         .  1..     .     . 
o =1.00211,  to  express  their  weights  in  terms 

of  that  of  the  ten-gram  piece  (10')  • 

What  we  actually  do,  assuming  (10')  is  found  to  have  an  abso* 
lute  value  of  9.99970  g.,  is  to  multiply  the  individual  values  of 
the  larger  pieces,  based  upon  (10')  as  a  standard,  by  0.99997, 
and  those  of  the  fractional  weights,  expressed  in  terms  of  the  ten- 
milligram  piece  (10'),  by"^'  ^^   ^ — Q- 999 5 4  y^  0.99997  =  1.00208. 

0.99908 

In  this  way,  we  finally  arrive  at  the  individual  values  of  the 
pieces,  both  large  and  small,  in  terms  of  the  absolute  standard. 
Errors  Due  to  Inequalities  in  Length  in  the  Beam  Arms.  In 
the  preceding  discussion,  it  has  mainly  been  assumed  that  the  two 
arms  of  the  beam  are  equal  in  length.  This  is  not  really  the  case. 
It  is  mechanically  impossible  to  insure  perfect  equality.  To  find 
the  relative  lengths  of  the  arms,  place  (corrected)  weights  of  the 
same  nominal  value  —  say,  50  grams  —  upon  each  pan,  and 
bring  the  balance  into  equilibrium  by  means  of  the  rider.  Inter- 
change the  weights  on  the  two  pans,  and  again  bring  the  balance 
into  equilibrium  by  means  of  the  rider.  Call  the  two  weights  W 
and  w,  and  let  /  and  r  respectively  denote  the  additional  weights 
required  for  equilibrium  on  the  left  and  right  sides.  Then,  on 
the  first  weighing,  w-{-l=W;  and,  on  the  second  weighing, 
W^w+r.  Let  L  and  R  respectively  denote  the  length  of  the 
left  and  right  arm.     Then  from  the  law  of  levers, 

L{w+l)  =  RW;  2ind  LW=R(w+r), 

Solving  each  of  these  equations  for  Wj  and  equating  the  results, 
we  find  that 


whence, 


INTRODUCTION  19 

Suppose,  for  example,  with  use  of  true  weights,  that  the  weigh- 
ings were  found  to  be : 

Left  Right 

(50)  =  (20)  +  (loO  +  (lo'O  +  (io"0  +0.13  mg. 

(2o)  +  (ioO  +  (io'0  +  (io"0  =  (5o)+o.i9mg. 

Here,  then,  /=  —0.00013,  r—  +0.00019,  and  2£^=  50  g.  Conse- 
quently if  in  the  above  expression  we  let  R—i,  we  have 

r     ^  \w-\-r 

L=\ — !— -=1.0000032 
w-\-l 

i.e.  L:  7^=1.0000032  : 1 

With  this  ratio,  —=1.0000032,  a  weight  w  on  the  left  pan  of 
R 

the  balance  will  be  equivalent  to  a  weight  ic^X  1.0000032  on  the 
right  pan.  Hence,  if  an  object  on  the  left  pan  balances  the 
weight  50  g.  on  the  right  pan,  the  weight  of  the  object  is 
50.0000X1.0000032=50.0016  g.  —  an  error  of  only  0.003  P^^^ 
cent.  There  is  therefore  no  need  to  apply  the  correction.  Each 
balance  has  its  own  constant  L:  Riot  sl  given  load ;  the  numer- 
ical value  of  the  ratio  varies  with  the  different  loads. 

Most  analytical  balances  do  not  require  a  correction  on  ac- 
count of  inequalities  in  the  arms ;  the  lengths  are  usually  suffi- 
ciently exact.  Anyhow,  if  the  weights  are  always  placed,  say, 
on  the  right  pan,  such  a  correction  is  unnecessary  in  ordinary 
analytical  work,  because  the  weights  observed  are  proportional 
to  the  true  weights,  and  the  ratios  obtained  are  not  affected. 

Errors  Due  to  Atmospheric  Buoyancy.  Two  bodies  equal 
in  weight  at  the  same  time  and  place  contain  the  same  mass  or 
quantity  of  matter  only  if  the  weighing  is  carried  out  in  a  vac- 
uum, or  if  the  bodies  have  the  same  volume.  The  latter  is 
hardly  ever  the  case. 

A  body  weighed  in  air  is  buoyed  up  by  a  pressure  equivalent 
to  the  weight  of  the  air  which  it  displaces.  Suppose  that  exactly 
10  grams  of  platinum  (sp.  gr.,  21.55)  are  weighed  in  air  with 


20  QUANTITATIVE  CHEMICAL  ANALYSIS 

brass  weights  (sp.  gr.,  8.4).  Then  0.45  cc.  of  air,  at  say  20°  and 
760  mm.,  i.e.  about  0.00054  g.,  are  displaced  by  the  platinum ; 
while  the  weight  of  air  displaced  by  10  grams  of  brass  is  0.00143 
g.  Hence,  the  weight  of  brass  which  will  be  required  to  counter- 
poise 10  grams  of  platinum  is  10+  (0.00143  —0.00054)  =  10.0009  g- 
The  arithmetic  of  the  above  calculation  is  summarized  in  the 
formula : 

Corrected  weight =2£;+w  /^— —  j 

in  which  w  represents  tne  apparent  weight  of  the  object ;  s,  the 
specific  gravity  of  the  object;  Si,  the  specific  gravity  of  the 
weights;  and  oj,  the  weight  of  a  cubic  centimeter  of  air  under 
the  conditions  prevailing  at  the  time  of  the  experiment. 

To  illustrate  the  effect  of  the  buoyancy  of  air  on  a  few  common 
substances,  when  weighed  with  brass  weights,  it  may  be  stated 
that  the  error  per  gram  of  substance  weighed  is  o.i  mg.  for  ferric 
oxide  (sp.  gr.  5.12);  0.3  mg.  for  Mg2P207  (sp.  gr.  2.40);  and 
0.4  mg.  for  sodium  chloride  (sp.  gr.  2.13). 

In  ordinary  analytical  operations  we  have  to  deal  with  differ- 
ences in  weight,  and  with  ratios,  not  with  absolute  weights. 
When  the  amount  of  a  precipitate  is  determined  from  the  dif- 
ference in  the  weight  of  an  empty  crucible  and  of  the  crucible 
plus  the  precipitate,  the  buoyancy  correction  is  not  needed  for 
precipitates  with  a  specific  gravity  near  that  of  the  substance 
undergoing  analysis.  If,  however,  the  specific  gravities  are 
widely  separated,  it  may  be  worth  while  to  correct  for  buoyancy. 
For  instance,  since  the  specific  gravities  of  iron  ore  and  barium 
sulphate  are  nearly  equal,  it  would  be  a  waste  of  time  to  correct 
for  buoyancy  in  determining  sulphur  in  an  ore  of  iron.  On  the 
other  hand,  in  standardizing  a  solution  of  silver  nitrate  by 
precipitating  silver  chloride  from  a  specific  weight  of  the  solution 
(use  of  weight  burettes),  the  buoyancy  of  air  may  cause  an  error 
of  0.1%. 

Summary.  The  foregoing  discussion  clearly  illustrates  the 
advisabiHty,  in  work  of  an  accurate  nature,  of  always  estimating 


INTRODUCTION  21 

the  effect  of  the  various  sources  of  error  on  the  final  result.  Any 
of  these  errors  can  then  be  neglected,  provided  it  is  sufficiently 
small  in  comparison  with  the  error  derived  from  other  sources. 
The  chief  sources  of  error  involved  in  weighing  are  those  due  to  : 
(i)  variations  in  the  rest-point  of  the  balance ;  (2)  inconsistent 
weights;  (3)  inequahties  in  length  in  the  beam  arms;  and 
(4)  atmospheric  buoyancy.  In  weighings  making  any  pretense 
to  accuracy,  the  following  points  should  be  noted : 

(i)  The  rest-point  of  the  unloaded  balance  should  be  deter- 
mined and  made  use  of  in  each  weighing. 

(2)  The  weights  should  be  caHbrated,  and  periodically  checked 
for  consistency. 

(3)  The  errors  due  to  inequahty  in  length  in  the  beam  arms 
may  be  neglected  in  ordinary  analytical  work. 

(4)  Although  the  correction  of  weighings  for  buoyancy  can 
almost  always  be  neglected  in  general  analytical  work,  owing 
to  the  much  larger  errors  associated  with  the  preparation  of 
precipitates  for  weighing,  it  is  advisable  in  weighing  bulky  glass 
apparatus  (potash  bulbs,  etc.),  to  use  a  similar  piece  of  appara- 
tus as  a  counterpoise.  This  will  serve  to  eliminate  any  errors 
due  to  variations  in  temperature,  pressure,  and  humidity  dur- 
ing the  course  of  the  determination. 

II.   PRECIPITATION 

Qualities  Desirable  in  Precipitates  Which  Are  to  be  Used  in 
Gravimetric  Determinations.  Precipitation  is  made  use  of  more 
often  than  any  other  means  for  the  separation  of  inorganic 
substances.  But,  in  carrying  out  such  separations,  precipitates 
should  conform  as  nearly  as  possible  to  the  following  ideal 
specifications:  (i)  The  precipitate  should  be  insoluble  in  the 
mother  liquid,  and  also  in  the  wash  Hquid  to  be  used;  (2)  it 
should  be  compact,  easy  to  filter  and  wash ;  (3)  it  should  be  a 
pure,  non-volatile  chemical  substance  of  known  percentage  com- 
position, or  it  should  yield  upon  ignition  a  pure,  non-volatile 
substance  of  known  composition.    The   last  condition  is  of 


22  QUANTITATIVE  CHEMICAL  ANALYSIS 

course  of  especial  importance,  if  the  precipitate  is  the  substance 
finally  to  be  weighed.  Moreover,  other  things  being  equal,  it 
is  conducive  to  accuracy  if  a  precipitate  can  be  obtained  which 
contains  a  low  percentage  of  the  substance  under  investigation 
(cf.  Part  V,  Problem  46). 

But  few  processes  satisfy  all  these  requirements,  and  in  the 
case  of  any  analytical  process  it  is  important  to  know  what 
conditions  favor  and  what  conditions  hinder  the  separation  and 
purification  of  a  given  precipitate.  There  are  a  few  general 
principles  of  such  wide  applicability  that  they  should  be  con- 
stantly borne  in  mind.     Their  discussion  follows. 

Colloidal  and  Fine-grained  Precipitates.  Finely  divided 
precipitates,  such  as  newly  precipitated  silver  chloride,  barium 
sulphate,  calcium  oxalate,  etc.,  are  particularly  liable  to  pass 
through  the  filter;  furthermore,  they  tend  in  large  measure  to 
stop  up  the  pores  of  the  filter,  and  thus  to  increase  the  time 
required  for  filtration  and  washing.  Hence,  the  analyst 
employs  various  artifices  in  order  to  enlarge  the  size  of  the 
particles. 

(i)  The  grain  size  can  frequently  be  increased  by  allowing 
the  fine  grains  which  originally  separate  to  digest  in  the  pre- 
cipitation liquid.  This  change  is  more  rapid  in  the  hot,  than 
in  the  cold,  mother  Hquid.  In  the  case  of  crystalHne  substances, 
it  often  happens  that  the  finer  grains,  which  (owing  to  differences 
in  surface  tension)  are  somewhat  more  soluble  than  the  coarser 
ones,  redissolve;  and  since  the  solution  is  then  supersaturated 
in  respect  to  the  coarser  grains,  these  are  augmented  in  size  by 
the  surface  deposition  of  the  material  furnished  by  the  finer 
grains. 

The  boiling  of  liquids  containing  colloidal  substances  fre- 
quently leads  to  the  flocculation  of  a  large  number  of  fine  par- 
ticles into  a  smaller  number  of  coarser  aggregates. 

(2)  Precipitates  produced  in  hot  solutions  are  often  coarser- 
grained  than  if  produced  in  cold  solutions.  From  what  has 
just  been  said,  the  reasons  for  this  fact  will  be  apparent. 


INTRODUCTION  23 

(3)  The  flocculation  of  a  precipitate  which  separates  in  a 
colloidal  condition  is  frequently  caused  by  the  salts  present  in 
the  mother  Hquid.  When  these  salts  have  been  almost  removed, 
during  the  washing,  the  colloidal  precipitate  is  apt  to  be  defloc- 
culated,  and  it  may  then  give  a  turbid  filtrate,  or  become  so 
slimy  as  to  be  almost  impermeable  to  the  wash  Hquid.  In  such 
cases  it  is  necessary  to  wash  the  precipitate  either  with  boiling 
water,  or,  better,  with  the  solution  of  an  electrolyte  which  will 
prevent  the  deflocculation  of  the  precipitate,  and  which  can  be 
easily  removed  by  drying  or  ignition.  Sometimes  dilute  acids 
can  be  used,  but  usually,  for  obvious  reasons,  we  have  to  de- 
pend upon  volatile  ammonium  salts. 

The  Contamination  of  Precipitates.  Finally,  it  should  be 
noted  that  the  finer  the  grain  of  the  precipitate,  the  greater  will 
be  the  quantity  of  contaminating  salts  likely  to  be  retained  by 
the  wet  precipitate.  The  salts  appear  to  be  retained  by  a  kind 
of  surface  attraction,  called  adsorption,^  and,  since  fine-grained 
precipitates  expose  a  larger  surface  of  separation  between  the 
solid  and  the  Uquid  phases,  and  also  because  of  their  compact- 
ness, the  fine  grained  precipitates  are  more  difficult  to  wash 
clean  than  those  of  coarser  texture.  Colloidal  gelatinous  pre- 
cipitates Hke  ferric  and  aluminum  hydroxides  are  in  an  ex- 
tremely fine  state  of  subdivision,  and,  in  consequence,  they  are 
most  difficult  to  wash  clean. 

In  addition  to  their  tendency  to  be  contaminated  by  the 
adsorption  of  salts,  precipitates  are  also  frequently  Uable  to 
contamination,  owing  to  the  formation  during  precipitation  of 
more  or  less  stable  insoluble  complexes  (and,  in  rare  cases, 
possibly,  to  the  carrying  down  of  foreign  substances  by  the 
precipitate  in  a  state  of  soHd  solution).  These  impurities,  in 
whatever  form  they  may  be  present,  cannot  be  completely  re- 
moved by  washing,  and  the  wash  water  will  frequently  fail  to 

*  But  see  "The  Contamination  of  Precipitates  in  Gravimetric  Analysis,"  G. 
McP.  Smith:  Journal  of  the  American  Chemical  Society ^  vol.  jp,  pp.  1152-73 
(1917). 


24  QUANTITATIVE  CHEMICAL  ANALYSIS 

show  any  indication  of  the  impurities  which  are  still  present  in 
the  precipitate. 

It  is  therefore  often  advisable  to  redissolve  the  precipitate, 
and  to  repeat  the  precipitation.  The  objectionable  impurity 
divides  itself  in  a  more  or  less  definite  concentration  ratio  be- 
tween the  precipitate  and  the  mother  liquid.  A  relatively  large 
amount  may  be  retained  by  the  precipitate  in  the  first  pre- 
cipitation, but  in  a  second  precipitation,  when  only  that  amount 
of  salt  retained  by  the  first  precipitate  is  in  solution,  the  divi- 
sion of  the  undesirable  substance  between  the  precipitate  and  the 
solution  in  the  given  concentration  ratio  means  that  a  much 
smaller  quantity  of  impurity  will  be  retained  by  the  second 
precipitate.  Repeated  precipitations  will,  in  general,  soon  carry 
the  amount  of  impurity  outside  the  range  of  the  balance ;  but, 
in  carrying  out  such  operations,  the  solubility  relations  of  the 
precipitate  itself  should  never  be  left  out  of  consideration. 

The  Theory  of  Precipitation.  Reversible  Reactions.  The 
reactions  which  are  made  use  of  in  analytical  chemistry  belong 
mostly  to  the  reversible  type.  Under  unfavorable  analytical 
conditions,  instead  of  running  to  completion,  such  reactions  may 
come  to  an  apparent  standstill,  owing  to  a  too  early  attainment  of 
equilibrium,  and  in  that  case  a  certain  quantity  of  the  substance 
involved  in  the  determination  is  likely  to  remain  unaccounted 
for.  This  is  true  in  many  reactions  involving  precipitation, 
oxidation  and  reduction,  and  even  neutralization.  Further- 
more, many  reactions  are  more  or  less  influenced  by  the  pres- 
ence of  certain  substances.  It  is  obvious  that  a  knowledge  of 
the  processes  which  take  place  in  such  cases  will  be  of  the  great- 
est service  to  the  analytical  chemist. 

Therefore,  it  is  of  the  first  importance  in  analytical  chemistry 
to  study  each  process  thoroughly  in  detail,  in  order  to  find  out  and 
understand  the  conditions  which  will  most  nearly  lead  to  the  com- 
pletion of  every  reaction  involved.  For  such  studies,  the  ionic 
theory  and  the  law  of  mass  action  are  indispensable  guides. 
It  is  taken  for  granted  that,  at  this  point,  the  student  is  already 


INTRODUCTION  25 

sufficiently  familiar  with  the  qualitative  conception  of  ioniza- 
tion. 

Degree  of  Ionization.  In  a  dilute  aqueous  salt  solution,  the 
greater/  and  by  far  the  most  active,  portion  of  the  salt  is  almost 
invariably  ionic.  But  with  acids  and  bases  there  is  a  wider 
range,  and  of  these  a  larger  number  are  less  highly  ionized ;  but 
even  here  the  ions  are  nearly  always  much  more  active  than  the 
non-ionized  molecules.  The  acids  and  bases  that  are  commonly 
called  "  strong "  are  highly  ionized,  i.e.  their  solutions  are 
especially  active  as  acids  or  bases  because  they  contain  high 
hydrogen,  or  hydroxide-ion,  concentrations. 

Composition  of  the  Ions. .  It  is  usual  to  assume  the  simplest 
possible  compositions  for  the  ions  formed  upon  the  dissociation 
of  any  given  electrolyte.  A  more  careful  study  of  the  subject, 
however,  shows  that  the  ionization  of  even  simple  electrolytes 
may  be  a  very  complicated  process.  It  is  known,  for  example, 
that  sulphuric  acid  contains  ions  of  the  formula  HS04~,  in  ad- 
dition to  SO4      ions,  and  that  phosphoric  acid  yields  ions  of 

the  formulas  H2P04~,  HPO4 — ,  and  PO4 .    All  of  these  are 

probably  more  or  less  highly  hydrated;  even  hydrogen  and 
hydroxide  ions  are  supposed  to  be  hydrated  in  aqueous  solu- 
tion. Further,  many  metallic  ions  show  a  decided  tendency  to 
exist  in  combination  with  certain  molecules  and  radicals,  as  OH2, 
OH,  NH3,  NH2,  CN,  C2O4,  PO4,  CI,  etc. ;  but  in  very  dilute 
solutions  these  complexes  are  apt  to  be  more  or  less  highly 
dissociated  into  their  constituents. 

The  Law  of  Mass  Action  as  Applied  to  Ionic  Equilibria.  In 
aqueous  solution,  acetic  acid  is  supposed  to  ionize  as  follows : 

HC2H3O2  :^H+  +C2H3O2- 

The  quantity  of  the  molecular  acid  that  is  ionized  per  unit  of 
time  in  a  given  volume  of  the  solution  is  proportional  to  the 
concentration  of  the  non-ionized  molecules,  ChcsHjOij  while  the 

*  Noteworthy  exceptions  are  mercuric  chloride  and  cyanide,  lead  acetate,  and 

a  few  others. 


26 


QUANTITATIVE  CHEMICAL  ANALYSIS 


quantity  of  the  molecular  acid  that  is  simultaneously  formed 
by  the  union  of  the  ions  depends  upon  the  frequency  of  the  en- 
counters of  the  two  kinds  of  ions,  which  in  turn  is  proportional 
to  the  product  of  their  concentrations,  Cn+XCcaHaOj-. 
The  speeds  of  the  respective  actions  will  therefore  be 

5i=Chc^30.XFi  and  52=Ch+  XQ^k),-  XF2, 

in  which  Fi  represents  the  intrinsic  tendency  of  HC2H3O2  to 
ionize,  and  F2  that  of  H+  and  C2H302~  to  combine. 

When  equal  amounts  of  material  are  being  transformed  each 
way,  i.e.  at  equilibrium,  Si^^S^y  and  therefore 

Chc2Hj02  X  /^i   =  Ch+  X  CcjHiOi"  X  Fti 
or  CH+XCc,H30.-_gi_j^  (j) 

ChchjO,         -^2 

— \  being  th«  ratio  of  two  constants,  is  constant ;  and  the  value, 
F2 

ky  of  this  ratio  of  the  affinities  driving  the  opposed  actions  is 
called  the  affinity  constant  of  the  reversible  reaction.  At  any 
given  temperature,  provided  the  solution  is  dilute,  the  numerical 
value  of  k  remains  the  same  no  matter  what  the  total  concentra- 
tion of  the  solution  may  be.^  In  the  case  of  acetic  acid,  for 
example,  the  following  figures  have  been  obtained,  at  18°,  from 
conductivity  determinations. 


Total  Moial 

Concentration 

OF  Acid 

Proportion 
Ionized 

MoLAL  Concentration 
OF  H  +  AND  of  Ac— 

(Ch+  and  CC2H30J-) 

MoLAL  Concentration 
OF  HAc 
(CHCJH302) 

I.OOO 

O.IOOO 

O.OIOO 

0.0041 
0.0130 
0.0407 

0.0041 

0.00130 

0.000407 

1. 000-0.0041 

O.IOOO-O.OOI30 

0.01000-0.000407 

*  When  data  such  as  the  following  are  applied  to  cases  of  soluble,  highly  ionized 
substances,  the  ^-values  so  obtained  for  any  given  compound  are  usually  far  from 
constant.  The  general  conclusions  arrived  at  through  the  application  of  such 
data  are,  however,  as  a  rule,  not  invalidated  by  this  fact. 


INTRODUCTION  27 

Substituting  these  figures  in  equation  (i)  above,  we  get : 

(0.0041)^  ^     (0.0013)2 

^^ -V- =  0.0000160 ;  ^^ =^  =  0.0000171 ; 

0.996  ^'   0.0987  ' 

and         V  -000407;  =QQQQQ^^2, 
0.00959 

It  is  seen  that,  although  the  third  solution  is  a  hundred  times 
more  dilute  than  the  first,  and  although  the  degree  of  ionization 
has  increased  tenfold,  the  value  of  k  is  the  same  in  both  cases. 

The  Common-Ion  Effect.  When,  through  the  presence  of 
two  substances  which  furnish  an  ion  in  common,  the  concentra- 
tions of  the  positive  and  negative  ions  of  an  ionogen  are  unequal, 
the  law  of  mass  action  still  holds. 

Let  us  imagine,  for  example,  that,  by  mixing  equal  volumes 
of  the  double-molal  solutions,  a  solution  is  obtained  which  is 
uni-molal  in  respect  to  acetic  acid  and  also  to  sodium  acetate. 
Let  us  further  suppose  that  the  equilibria  which  exist  in  the 
mixture  have  been  established  in  two  separate  stages,  as  follows : 
(i)  that  the  concentrations  of  each  undissociated  compound 
and  its  ions  have  changed  from  those  which  exist  in  a  double- 
molal,  to  those  which  exist  in  a  uni-molal  solution  of  that  com- 
pound; and  (2)  that  the  concentrations  of  all  the  substances 
present  have  changed  from  those  which  exist  in  the  separate 
uni-molal  solutions  of  the  compounds,  to  those  which  exist  in 
the  mixture  which  is  uni-molal  in  respect  to  each  compound. 
Let  us  now  consider  this  latter  stage  in  detail. 

In  uni-molal  solution,  sodium  acetate  is  0.53  ionized,  while 
acetic  acid  at  that  concentration  is  only  0.004  ionized.  Each 
compound  furnishes  acetate  ions,  and  the  acetate  ions  present  are 
all  available,  either  for  union  with  sodium  ions,  or  for  union  with 
hydrogen  ions.  Initially,  therefore,  in  the  case  of  the  sodium 
acetate,  we   have  ^-53X0.534^^^^  instead  of  ^-53X0.53^^. 

0.47  0.47 

but  the  two  expressions  are  so  nearly  identical  that  we  see 
at  a  glance  that  the  ionic  equiHbrium  of  the  salt  will  not  be 


28  QUANTITATIVE  CHEMICAL  ANALYSIS 

affected  appreciably  by  the  presence  of  the  acid.  In  the  case 
of  the  acetic  acid,  however,  we  have  the  initial  relationship, 
0.004X0.534^  133  k,  instead  of  °-'^4X 0.004^;^     Since,  at  equi- 

0.996  0.996 

librium,  the  fraction  -^-^^ — CiKsOt    j-e^iains  constant,  and  since, 

owing  to  the  low  H+-ion  concentration,  CHC2H3O2J  cannot  be  in- 
creased appreciably,  nor  CcjHaOj-  be  appreciably  diminished, 
by  the  formation  of  the  molecular  acid,  it  follows  that  the  value 
of  Ch+  must  be  lowered  to  about  xlir  of  its  initial  magnitude. 
That  is  to  say,  the  sodium  acetate  in  this  solution  diminishes 
the  hydrogen-ion  concentration  from  0.004  to  about  0.00003. 

The  student  should  especially  note  that  the  concentration  of 
a  given  ion  can  be  lowered  in  this  way  to  a  value  approximating 
zero  only  when  that  ion  unites  with  an  ion  added  to  form  a  substance 
which  is  insoluble  or  which  by  nature  has  only  a  very  slight  tendency 
to  dissociate.  We  might  add  sodium  chloride  in  the  hope  of 
repressing  the  ionization  of  hydrochloric  acid;  but,  since  both 
compounds  ionize  highly,  we  should  obtain  no  appreciable 
effect.  If,  however,  we  add  sodium  acetate  in  excess  to  hydro- 
chloric acid,  we  can  obtain  a  solution  which  is  as  weakly  acid 
as  the  one  discussed  above.^ 

The  Solubility  Product.  One  of  the  commonest  and  most 
interesting  appHcations  of  the  law  of  mass  action  is  met  with 
in  connection  with  the  precipitation  and  solution  of  relatively 
insoluble  salts. 

Every  substance  possesses,  when  immersed  in  a  liquid,  a 
certain  solution-tension,  by  which  is  meant  an  expansive  force 
which  tends  to  drive  particles  of  the  substance  outward  into 
the  liquid.  These  particles  move  in  every  direction,  and  conse- 
quently some  of  them  return  to  the  solid  and  reattach  them- 

1  For  example,  i  mol  of  HCI+2  mols  of  NaC2H302,  in  a  volume  of  i  liter,  give 
a  solution  which  is  uni-molal  in  respect  to  acetic  acid,  to  sodium  acetate,  and  to 
sodium  chloride.  The  hydrogen-ion  concentration  of  this  solution  would  also 
approximate  0.00003. 


INTRODUCTION  29 

selves  to  it.  This  occurs  the  more  and  more  frequently,  as  the 
concentration  of  the  particles  in  the  Hquid  increases,  until, 
finally,  a  stage  is  reached  at  which  the  number  of  particles 
leaving  the  solid  per  unit  of  time  is  equal  to  the  number  deposited 
upon  it.  When  the  entire  liquid  is  equally  charged  with  dissolved 
particles,  the  liquid  immediately  surrounding  the  soHd  will  lose 
none  by  diffusion,  and  a  condition  of  equilibrium  will  be  estab- 
lished. At  a  constant  temperature,  the  quantity  of  dissolved 
solute  will  remain  thereafter  unchanged,  no  matter  how  long  the 
materials  are  left  in  contact.  It  is  at  this  point  that  the  solution 
is  said  to  be  saturated  with  the  dissolved  substance. 

In  the  case  of  silver  bromate  in  water,  we  have  the  following 
scheme  of  equilibria : 

AgBrOa  X  AgBrOs  X  Ag++Br03-. 

(solid)  (dissolved) 

The  solid  AgBrOs  molecules  tend  to  enter  the  solution,  while  at 
the  same  time  dissolved  AgBrOs  molecules  tend  to  come  out  of 
solution,  and  the  solution  is  saturated  when  these  tendencies 
produce  equal  effects.  The  ions  themselves  (and  any  foreign 
materials  present)  are  not  supposed  to  take  any  direct  part  in 
the  equilibrium  which  controls  solubility.  That  is,  in  solutions 
saturated  at  a  given  temperature  by  a  given  solute,  the  concentra- 
tion of  the  non-ionized  molecules  will  he  constant  no  matter  what 
other  substances  may  be  present,  provided  that  the  quantities  of  all  the 
dissolved  substances  are  not  sufficient  to  alter  the  nature  of  the  solvent. 
The  total  solubility  of  an  ionogen,  as  we  ordinarily  use  the  term, 
is  made  up  of  a  molecular  and  an  ionic  part.  The  quantity  of  the 
latter  does  not  remain  constant  when  a  foreign  substance  giving 
a  common  ion  is  added  to  the  solution.  In  a  solution  of  silver 
bromate,  for  example,  we  have  the  mathematical  relationship : 

(^^J4(?I^=^,  or  (Ag+)X(Br03-)  =  ;^X(AgBr03). 
(AgBrOs) 

But,  since,  in  the  special  case  of  a  solution  which  is  saturated  with 
the  salt  at  a  given  temperature,  the  concentration  of  the  non- 


30 


QUANTITATIVE  CHEMICAL  ANALYSIS 


ionized  molecules,  (AgBrOa),  remains  constant,  it  follows  that 
the  product,  y^xCAgBrOs),  also  remains  constant,  or  that  in  a 
saturated  solution  of  a  given  slightly  soluble  ionogen  the  product 
of  the  concentrations  of  its  ions  is  constant.  This  product  is  called 
the  solubility  product,  because  the  two  separate  values  jointly 
determine  the  magnitude  of  the  total  solubility  of  the  ionogen. 
The  concentration  of  the  non-ionized  molecules  cannot  be 
diminished,  but  the  ionic  part  of  the  solute  may  become  vanish- 
ingly  small  if  the  concentration  of  the  common  ion  is  made  great 
as  compared  with  that  of  the  other  ion  of  the  solute.  The 
relationships  which  exist  in  the  case  of  silver  bromate  are  illus- 
trated in  the  following  table,  where  it  will  be  seen  that,  in  this 
instance,  the  experimental  values  agree  remarkably  well  with 
the  calculated  ones. 


Solubility  of  AgBrOs  in  Mols  per  Liter 


MoLS  PER  Liter  of 

Solubility  Found 

Solubility  Calc. 

Common-ion  Salt 
Added 

Addition  of  Silver 
Nitrate 

Addition  of  Potassium 
Bromate 

(Addition  of  either 
Salt) 

0.00810 
0.00510 
0.00216 

0.00810 
0.00519 
0.00227 

O 

0.00850 

0.0346 

0.00504 
0.00206 

The  theory  of  the  precipitation  and  solution  of  slightly  soluble 

ionogens  may  be  summed  up  as  follows :  ^ 

^  That  is,  of  uni-univalent  ionogens.  In  other  cases,  the  solubility  pioduct 
would  often  contain  ion-concentrations  raised  to  the  second,  third,  etc.,  powers; 
but  in  reality  the  question  is  very  much  complicated  by  interfering  reactions. 
Thus,  in  the  case  of  PbCla,  if  NaCl  is  added  to  the  saturated  solution,  some  PbCl2 
will  be  precipitated  in  accordance  with  the  theory;  but  the  addition  of  Pb(N03)2 
actually  increases  the  solubility  of  the  PbCU.  This  is  probably  because  of  the 
formation  of  complexes,  or  of  intermediate  ions,  such  as  PbCl+,  or  of  both,  whereby 
the  addition  of  the  salt  giving  the  common  bivalent  ion  may  not  only  fail  to  increase 
the  concentration  of  the  bivalent  ion,  but  may  even  lower  that  of  the  univalent  ion. 
At  any  rate,  enough  is  known  to  indicate  that  the  theory  may  not  be  so  much  at 
fault  as  we  ourselves,  in  our  lack  of  methods  for  finding  out  just  what  ions  and 
complexes  are  present  in  such  solutions,  and  in  what  quantities. 


INTRODUCTION  31 

Whenever  the  ion-concentration  product  of  any  pair  of  ions  in  a 
solution  is  made  to  exceed  in  value  the  solubility  product  of  the  com- 
pound formed  hy  their  union,  the  latter  will  he  precipitated  until 
the  ion-concentration  product  has  been  reduced  to  the  solubility- 
product  value.  And  conversely,  whenever  the  ion-concentration 
product  of  any  pair  of  ions  is  made  less  than  the  solubility-product 
value,  the  compound  formed  by  their  union  will  be  dissolved  by  the 
solution  J  if  furnished  in  excess,  until  the  solubility-product  value 
has  been  regained. 

III.  FILTRATION  AND  THE  WASHING  OF  PRECIPITATES 

The  purpose  of  filtration  is  to  separate  a  solid  from  a  liquid 
in  which  it  is  suspended.  This  is  effected  by  causing  the  liquid 
to  pass  through  a  porous  medium  compact  enough  to  retain  the 
solid.  The  most  important  media  in  use  are  filter  paper,  as- 
bestos pulp,  and  platinum  sponge. 

The  Selection  and  Use  of  Paper  Filters.  Three  qualities 
which  are  desirable  in  a  filter  are :  (i)  porosity,  to  insure  rapid 
filtration;  (2)  sufficient  compactness,  to  insure  complete  re- 
tention of  the  precipitate ;  and  (3)  low  amount  of  ash.  Filters 
for  quantitative  work,  readily  obtainable  on  the  market,  are 
such  as  have  been  treated  with  hydrochloric  and  hydrofluoric 
acids.  Upon  incineration,  they  leave  a  small,  definitely  known 
weight  of  ash.^ 

Rapid  (porous)  filters  should  be  used  for  all  precipitates  which 
do  not  pass  through  them,  and  slow  (compact)  papers  should 
not  be  used  unless  necessary.  A  great  amount  of  time  is  con- 
sumed, often  wasted,  in  the  filtration  and  washing  of  precipitates. 

The  size  of  the  filter  should  be  determined,  not  by  the  volume 
of  liquid  to  be  filtered,  but  by  that  of  the  precipitate ;  the  latter 
should  not  more  than  half  fill  the  paper,  but  too  large  a  paper 
leads  to  a  waste  of  time  in  washing.  The  filter,  as  well  as  the 
precipitate,  retains  certain  salts  very  tenaciously. 

*  For  example  each  9  cm.  filter  of  a  certain  well-known  brand  leaves  on  ignition 
an  ash  weighing  about  o.ii  mg. 


32  QUANTITATIVE  CHEMICAL  ANALYSIS 

Funnels  should  be  provided  with  narrow  stems  of  even  bore, 
about  eight  inches  long,  and  should  have  an  angle  of  60°,  so  that 
a  twice  folded  filter  when  opened  up  will  closely  fit  its  walls; 
the  top  edge  of  the  filter  should  always  be  well  below  the  rim  of 
the  funnel. 

Upon  being  placed  in  the  funnel,  the  paper  should  be  wetted, 
and  carefully  bedded  against  the  funnel  walls,  so  that,  when 
water  is  poured  in,  the  stem  of  the  funnel  will  be  filled,  without 
the  entrance  of  air.^  If  properly  bedded  in  a  clean  funnel,  water 
will  rapidly  pass  through,  and  the  paper  will  be  at  its  best  for 
service  as  a  filter ;  the  filtration  proceeds  under  a  pressure  equiva- 
lent to  the  weight  of  a  column  of  water  of  the  same  diameter  as 
the  bore  of  the  stem  and  of  a  height  equal  to  the  length  of  the 
stem  plus  the  depth  of  liquid  in  the  filter. 

If  suction  is  to  be  appHed,  the  paper  should  preferably  be  com- 
pact, and  its  apex  should  be  supported  by  a  perforated  cone  of 
platinum  (or  of  palau),  placed  beneath  the  filter  in  the  funnel. 
In  general,  if  accuracy  is  aimed  at,  one  should  hesitate  to  use 
suction  with  paper  filters. 

After  the  precipitate  has  been  allowed  to  settle,  the  clear 
liquid  should  be  transferred  to  the  filter  by  means  of  a  glass  rod, 
along  which  it  should  be  made  to  flow  towards  the  side,  not  the 
center  of  the  paper.  The  outlet  of  the  funnel  should  rest  against 
the  inner  side  of  the  receiving  vessel,  to  enable  the  liquid  to  run 
down  quietly,  without  danger  of  loss  by  splashing. 

Before  beginning  to  wash  a  precipitate,  the  receiving  vessel 
for  the  filtrate  should  always  be  replaced  by  a  clean  beaker; 
although  the  precipitate  may  have  shown  no  tendency  to  pass 
through  the  filter,  it  is  often  possible  that  it  may  be  taken  into 
colloidal  solution  by  the  wash  liquid. 

In  general,  the  precipitate  can  be  most  efficiently  washed  by 
decantation,  i.e.  by  the  addition  of  successive  portions  of  wash 

1  If  this  fails  to  take  place,  either  the  stem  of  the  funnel  is  too  wide,  or  it  is 
not  free  from  grease.  The  latter  can  be  removed  by  means  of  warm  cleaning 
solution. 


INTRODUCTION  33 

liquid,  followed  by  settling  and  the  transfer  of  the  clear  liquid 
to  the  filter.  Finally  the  precipitate  itself  may  be  transferred, 
and  the  washing  completed  upon  the  filter. 

It  will  always  be  found  that  small  portions  of  the  precipitate 
adhere  to  the  walls  and  bottom  of  the  containing  vessel.  These 
can  be  rubbed  loose  by  means  of  a  so-called  "  policeman y^^  —  a 
piece  of  soft  rubber  tubing  tightly  fitted  on  the  end  of  a  glass 
rod.  Pieces  of  rubber  tubing  with  closed  ends  are  sold  for  the 
purpose.  These  rubber-tipped  rods  should  be  used  only  for  the 
above-mentioned  purpose;  they  should  never  be  used  as  stirring 
rods  or  be  allowed  to  stand  in  analytical  solutions. 

The  washing  of  precipitates  should  always  be  completed 
promptly  after  filtration;  if  allowed  to  stand,  they  are  apt  to 
dry  out  and  crack,  with  the  formation  of  channels  for  the  free 
passage  of  wash  liquid. 

After  filtering  off  a  precipitate,  the  filtrate  should  always  be 
tested  to  insure  complete  precipitation,  and,  after  several  wash- 
ings have  been  made,  the  last  washings  should  be  tested ;  only  a 
few  drops  of  the  latter  should  be  taken  at  first  if  the  filtrate  con- 
tains substances  to  be  subsequently  determined,  but  near  the 
end  of  the  washing  several  cubic  centimeters  should  be  used. 
The  necessity  of  making  these  tests  cannot  be  too  strongly  im- 
pressed upon  the  student;  no  exception  should  be  made. 

Wash  Bottles.  Wash  bottles  for  use  in  quantitative  work 
usually  consist  of  rubber-stoppered  flasks  of  250-500  cc.  ca- 
pacity, of  resistance  glass,  with  tubes  smoothly  bent,  and  with  a 
jet  made  flexible  through  a  joint  with  the  outlet  tube,  by  means 
of  a  short  piece  of  black  rubber  tubing.  The  jet  should  deliver 
a  thin  even  stream  of  liquid.  For  use  with  hot  water,  the  neck  of 
the  flask  may  be  wrapped  with  asbestos  twine,  or  other  material. 

Gooch's  Filtration  Crucible.  In  1878,  F.  A.  Gooch  proposed, 
for  the  separation  of  certain  precipitates  by  filtration  with  suc- 
tion, the  use  of  a  mat  of  asbestos  bedded  on  the  perforated  bottom 
of  a  crucible,  thus  rendering  it  possible  to  wash,  dry,  and  weigh 
the  precipitate  in  the  crucible. 


34  QUANTITATIVE  CHEMICAL  ANALYSIS 

Preparation  of  the  Asbestos.  Of  the  several  varieties  of 
asbestos  on  the  market,  the  long-fiber,  silky,  chrysoHte  asbestos 
is  best  suited  for  this  purpose.  The  asbestos  should  be  rubbed 
roughly  over  the  surface  of  a  lo-mesh  brass  sieve,  inverted  on  a 
piece  of  paper,  until  a  sufficient  amount  has  passed  through. 
This  is  shaken  up  with  water,  allowed  for  the  most  part  to  settle, 
and  the  very  fine  particles  poured  off  with  the  water.  The  pulp 
is  then  digested  for  one  hour  on  the  steam  bath  with  concentrated 
hydrochloric  acid,  in  a  covered  porcelain  dish.  At  the  end  of 
this  operation,  water  is  added,  and  the  liquid  is  poured  off  through 
a  funnel  provided  with  a  platinum  filtration  cone ;  the  asbestos 
being  washed  with  hot  water,  at  first  by  decantation,  and  finally 
in  the  funnel,  with  gentle  suction,  until  the  washings  give  no 
opalescence  with  silver  nitrate  solution.  The  washed  asbestos 
is  finally  mixed  with  distilled  water,  and  preserved  in  a  glass- 
stoppered  bottle. 

Preparation  of  the  Gooch  Filter,  The  Gooch  crucible  is  con- 
nected with  the  suction  apparatus  by  means  of  an  adapter  — 
a  glass  cyhnder  5-6  cm.  long,  of  about  3  cm.  inside  diameter, 
which  is  drawn  out  at  the  bottom  into  a  tube  of  about  the  same 
length,  for  passage  through  the  rubber  stopper  of  the  suction 
flask.  A  4-cm.  length  of  band  tubing  is  stretched  over  the 
mouth  of  the  adapter,  so  that  about  half  of  it  projects  above  the 
top,  and  the  crucible  is  pressed  into  this  projecting  end  of  the 
rubber  tube. 

With  the  connections  all  well  made,  suction  is  gently  applied, 
and  asbestos  suspension  is  gradually  added  until  a  smooth  mat, 
not  over  1.5  mm.  in  thickness,  is  formed;  this  is  then  covered 
with  a  thin  perforated  porcelain  disk  (filter  plate),  and  enough 
asbestos  added  to  barely  cover  the  disk.  With  increased  suc- 
tion, water  is  next  filtered  through  the  crucible  until,  when 
held  before  a  bright  light,  no  trace  of  suspended  asbsetos  can 
be  seen  in  the  washings.  Usually,  250-500  cc.  of  water  will 
suffice  to  render  the  mat  stable  upon  further  washing.  The 
filter  plate  is  used  to  protect  the  asbestos  mat  during  filtration 


INTRODUCTION  35 

and  washing ;  if  the  perforations  of  the  crucible  are  rather  large, 
it  is  advisable  also  to  use  a  filter  plate  beneath  the  mat. 

After  the  preparation  of  a  satisfactory  mat,  the  crucible, 
placed  in  a  little  beaker,  is  dried  for  an  hour  at  120-130°,  in  an 
oven ;  it  is  then  allowed  to  cool  in  a  desiccator,  and  weighed. 
This  process  is  repeatedj  until  there  is  no  further  loss  in 
weight. 

Use  of  the  Gooch  Filter,  With  the  weighed  crucible  fitted 
in  the  adapter,  and  with  gentle  suction  already  applied,  the 
liquid  should  be  poured  from  the  precipitate  into  the  crucible ; 
the  precipitate  may  then  be  washed  by  decantation,  trans- 
ferred to  the  crucible,  and  the  washing  there  completed.  The 
rod  used  to  transfer  the  liquid  should  reach  nearly  to  the  bot- 
tom of  the  crucible,  in  order  not  to  stir  up  the  asbestos,  and  the 
filtrate  should  be  carefully  examined  for  suspended  particles  of 
asbestos,  and  repassed  through  the  filter  if  any  are  visible.  The 
crucible  with  its  contents  is  finally  dried  to  constant  weight,  as 
already  described.  The  increase  represents  the  weight  of  the 
precipitate. 

The  convenience  and  efficiency  with  which  precipitates  can 
be  prepared  for  the  balance  (with  the  repeated  use  of  the  same 
filter),  leads  to  the  use  of  crucibles  of  this  type  with  practically 
all  coarse-grained,  crystalline  precipitates ;  colloidal,  gelatinous 
precipitates  tend  to  stop  up  the  pores  of  the  filter. 

With  Gooch  crucibles  of  platinum,  set  into  closely  fitting, 
shallow  caps  of  platinum,  the  precipitates  may  readily  be  ig- 
nited at  a  high  temperature ;  but  in  this  case  it  is  better  to  pack 
the  crucible  with  a  felt  of  platinum  sponge.^  In  this  form,  the 
crucible  is  known  as  a  Munroe  crucible,  and,  if  properly  pre- 
pared, the  platinum  felt  is  far  more  suitable  for  fine-grained 
precipitates  than  that  of  asbestos.  Gold  crucibles  provided  with 
platinum  filters  are  cheaper,  but  should  not  be  heated  to  a  high 
temperature. 

*  For  the  preparation  of  such  filters,  see  Journal  of  the  American  Chemical 
Society,  vol.  31,  p.  456. 


36  QUANTITATIVE  CHEMICAL  ANALYSIS 

Sources  of  Error.  Since  asbestos  may  absorb  appreciable 
amounts  of  alkali,  not  removed  by  washing,  solutions  contain- 
ing fixed  alkalies  should  as  a  rule  not  be  filtered  through  as- 
bestos. Also,  asbestos  is  slightly  attacked  by  water  and  dilute 
acids;  after  the  treatment  indicated  above,  however,  there  is 
not  much  danger  from  this  source.  •  Munroe  crucibles,  on  the 
whole,  give  less  chance  for  error. 

The  Theory  of  Washing  Precipitates.  The  theory  of  wash- 
ing precipitates  should  include  the  consideration  of  several 
factors,  among  which  may  be  mentioned  the  phenomena  of 
adsorption  (in  which  the  filter  also  takes  part),  and  the  tend- 
ency of  precipitates  to  enter  the  wash  liquid  in  colloidal  form. 
But  these  factors  have  been  discussed  under  precipitation. 
Aside  from  these  considerations,  there  is  the  important  ques- 
tion concerning  the  most  effective  method  of  washing  precipitates. 

Let  us  suppose  that  a  precipitate  is  to  be  washed  by  decanta- 
tion,  and,  for  the  sake  of  simplicity,  that  neither  it  nor  the  filter 
exercises  any  physical  or  chemical  action  on  the  dissolved  salts. 
The  mother  liquid  has  been  decanted  as  far  as  possible  into  the 
filter,  and  the  latter  allowed  to  drain. 

Let  V  cc.  be  the  total  volume  of  solution  which  remains  in 
contact  with  the  precipitate  and  filter,  and  V  cc.  the  volume  of 
wash  water  added  each  time ;  and  assume  that  the  latter  mixes 
uniformly  with  the  liquid  adhering  to  the  precipitate  and  filter. 
Then,  upon  the  addition  of  V  cc.  of  water,  the  total  volume  of 
liquid  is  {V-{-v)  cc.  Further,  let  Co  be  the  concentration  in 
grams  per  cubic  centimeter  of  the  salts  in  the  original  so- 
lution; then  the  quantity  contained  in  the  v  cc.  left  in  con- 
tact with  the  precipitate  and  filter  is  vCfi  grams.  By  the 
addition  of   V  cc.  of  water,  the  concentration  is  reduced  to 

Ci=  — — Co,  and  if  this  liquid  is  removed  until  only  v  cc.  are 
V-\-v 

left,  the  quantity  of  undesirable  salts  present  is  reduced  to 

7) 

vCi  =  77 — .^Co  gm.    A  second  addition  of  V  cc.  of  water  gives  the 


INTRODUCTION  37 

concentration,  C2=  77 — Ci=(— — )  Co,  and  the  quantity  of 

V-\-v         \V-\-vJ 

salts  in  the  v  cc.  left  on  draining  is  now 

or,  after  n  washings,  the  quantity  of  undesirable  salts  has  been 
diminished  to  the  value, 

Y 


^-(ir^j^Cog-- 


This  formula  expresses  mathematically  the  self-evident  fact 
that,  for  a  given  number  of  washings,  the  quantity  of  undesir- 
able salts  left  behind  will  be  the  smaller,  the  more  completely 
the  precipitate  and  filter  are  drained,  and  the  greater  the  volume 
of  the  wash  water  that  is  added  each  time.  The  formula  enables 
us,  however,  to  answer  a  less  simple  question;  viz..  What 
is  the  most  efficient  method  of  washing  a  precipitate  with 
a  given  amount  of  wash  liquid  ?  Suppose,  for  example,  we  wish 
to  use  150  cc.  of  wash  liquid:  is  it  better  to  wash  six  times 
with  25  cc.  portions, or  to  wash  10  times  with  15  cc.  portions? 
Let  us  assume  that  Co=o.  i  g.  per  cubic  centimeter,  and  that 
i;=5  cc. ;  then,  in  the  two  cases,  the  quantities  of  undesirable 
salts  left  behind  will  be  (^V)^Xo.5=o.cxxx)io7  g.,  and  (^)^^Xo.5 
=  0.00000047  g.,  respectively.  Disregarding  adsorption,  which 
greatly  decreases  the  efficiency  of  washing,  ten  washings  with 
15  cc.  portions  are  23  times  as  efficient  as  six  washings  with 
25  cc.  portions.  Both  methods  of  procedure  will  require  ap- 
proximately equal  intervals  of  time,  since,  in  either  case,  150  cc. 
of  Hquid  must  run  through  the  filter.  It  is  much  better  to  wash 
a  precipitate  many  times  with  small  portions  of  liquid,  than  a  few 
times  with  larger  portions.  Each  portion  of  wash  liquid  should 
he  removed  as  far  as  possible  by  decantation  and  drainage,  before 
the  addition  of  a  fresh  portion. 

Another  factor  to  be  considered  is  the  temperature  of  the 


38  QUANTITATIVE  CHEMICAL  ANALYSIS 

solution  to  be  filtered.  Since  the  rate  of  flow  through  a  filter 
depends  largely  upon  the  viscosity  of  the  Hquid,  and  since  the 
viscosity  of  water  at  ioo°  is  only  one  sixth  that  at  o°,  it  is  well 
to  filter  and  wash  at  a  high  temperature,  unless  there  is  good 
reason  to  the  contrary. 

Finally,  in  washing  a  precipitate  on  a  paper  filter,  great  care 
must  be  taken  to  wash  the  filter  itself.  Soluble  salts  are  tena- 
ciously held  back  at  the  upper  edges  of  the  paper,  and  therefore 
this  part  of  the  filter  should  receive  especial  attention.  //  is 
best  to  fill  the  filter  each  time,  and,  before  refilling,  to  allow  it  to 
drain  completely.  Filters  should  be  selected  which  are  not  too 
large. 

IV.  THE  DRYING  AND  IGNITION  OF  PRECIPITATES 

Drying  Ovens.  There  are  on  the  market  many  types  of 
drying  ovens,  heated  by  gas,  by  steam  pipes,  or  by  electricity, 
in  which  the  temperature  may  be  more  or  less  accurately  con- 
trolled. The  oven  consists  essentially  of  a  drying  chamber, 
through  which  there  is  provided  a  slow  circulation  of  hot  air. 

A  precipitate  is  dried  on  the  filter  by  placing  the  funnel  con- 
taining both  in  a  drying  oven,  at  90-100°,  and  leaving  it  there 
for  a  sufficient  time.  The  funnel  should  be  covered  with  a 
sheet  of  common  filter  paper,  fastened  in  place  by  crimping  its 
edges  over  those  of  the  funnel.  If  the  precipitate  is  suitable 
for  weighing  without  ignition,  e.g.  silver  chloride  in  a  Gooch 
crucible,  it  should  be  dried  to  constant  weight  at  a  temperature 
considerably  above  the  boiling  point  of  water  (in  this  case,  at 
120-130°),  in  order  to  remove  the  last  traces  of  moisture  from 
the  filter  as  well  as  from  the  precipitate.  (Before  it  is  used,  the 
packed  Gooch  crucible  should,  of  course,  be  dried  to  constant 
weight  at  the  same  temperature.) 

Many  precipitates  may,  under  proper  precautions,  be  ignited 
without  previous  drying  in  an  oven.  But  if  such  precipitates 
can  be  dried  over  night,  or  otherwise,  without  loss  of  time  to 
the  analyst,  it  is  well  to  submit  them  to  this  process. 


INTRODUCTION  39 

The  precipitate,  folded  within  the  filter,  and  placed  at  the 
bottom  of  the  inclined  open  crucible,  is  heated  gently,  to  avoid 
the  sudden  escape  of  steam.  The  drying  is  most  safely  ac- 
complished by  alternately  applying  and  withdrawing  the  flame, 
the  heat  being  applied  to  the  lower  side  of  the  crucible,  near  its 
mouth.  When  the  bundle  is  dry,  the  crucible  is  heated  at  the 
same  point,  but  with  a  stationary  flame,  the  temperature  being 
gradually  increased  to  the  charring  point  of  the  paper;  the 
volatile  products  should  pass  off  without  burning,  since  that 
would  give  rise  to  draughts  within  the  crucible,  with  consequent 
danger  of  loss.  After  the  dry  distillation  of  the  paper,  in  burn- 
ing off  the  carbon,  the  flame  should  be  appUed  to  the  bottom 
of  the  crucible,  since  this  will  insure  a  gentle  circulation  of  air 
within  the  inclined  vessel.  The  crucible  may  finally  be  heated 
to  redness  until  the  ignition  is  complete. 

Some  precipitates  are  reduced  or  otherwise  affected  by  the 
hot  carbon  or  reducing  gases  from  the  filter  paper;  e.g.  silver 
chloride,  lead  sulphate,  etc.  are  reduced  to  metal  Since,  how- 
ever, these  metals  are  volatile  only  at  very  high  temperatures, 
there  is  no  loss  in  their  case,  and  the  metal  can  easily  be  con- 
verted into  the  original  compound.  In  such  cases,  it  is  advisable 
to  separate  the  precipitate  as  far  as  possible  from  the  filter,  and 
then  to  ignite  the  latter.  The  small  quantity  of  reduced  metal 
is  moistened  with  a  few  drops  of  nitric  acid,  and  the  resulting 
nitrate  converted  into  silver  chloride  with  hydrochloric  acid, 
or  into  lead  sulphate  with  sulphuric  acid,  and  the  excess  of  acid 
expelled  by  cautiously  heating  the  crucible.  The  bulk  of  the 
precipitate  is  then  added,  and  the  whole  ignited. 

Only  in  special  cases  may  precipitates  safely  be  heated  over 
the  blast  lamp,  and  even  then  it  is  often  preferable  to  use  a  large 
Meker  burner. 

Desiccators.  In  order  to  be  weighed  with  accuracy,  a  freshly 
ignited  object  must  first  be  allowed  to  cool ;  but  during  this  pro- 
cess it  should  be  protected  from  the  action  of  moisture,  carbon 
dioxide,  etc.    It  is  therefore  necessary  to  use  a  desiccator,  —  a 


40  QUANTITATIVE  CHEMICAL  ANALYSIS 

special  form  of  glass  vessel,  rendered  air-tight  by  means  of  ground- 
glass  contact  surfaces  which  are  thinly  coated  with  grease.^ 

For  general  analytical  work,  desiccators  are  usually  charged 
with  coarse  granules  of  anhydrous  calcium  chloride,  though 
sulphuric  acid  (spread  out  over  fragments  of  pumice,  to  increase 
the  active  surface),  soUd  potassium  hydroxide,  or  quick-lime  is 
sometimes  used.  Supported  above  the  drying  agent,  there  is  a 
porcelain  plate  which  is  provided  with  holes  of  a  size  suitable 
for  the  reception  of  crucibles.  Desiccators  should  be  opened 
only  when  necessary,  and  promptly  closed ;  the  charge  will  then 
retain  its  efficiency  for  many  months. 

Crucibles.  The  most  commonly  used  crucibles  are  of  high 
grade  porcelain.  They  withstand  very  high  temperatures  with- 
out appreciable  change  of  weight,  and  are  comparatively  cheap. 
They  are  not  suitable  for  fusions  because  most  fluxes,  particu- 
larly basic  ones,  attack  the  glaze  as  well  as  the  porcelain.  Even 
in  the  ignition  of  ordinary  precipitates,  in  spite  of  careful  wash- 
ing, traces  of  fusible  material  usually  remain  present,  and  these 
in  time  roughen  the  glaze  and  destroy  the  usefulness  of  the 
crucible. 

Crucibles  more  or  less  suitable  for  the  ignition  of  precipitates 
are  also  made  of  alundum,  and  of  fused  siUca. 

Crucibles  of  platinum  are  often  very  desirable  for  ignitions ; 
and  for  many  fusions  they  are  essential.  Platinum  melts  at 
about  1770°  and  does  not  soften  enough  to  preclude  its  use  at 
temperatures  slightly  below  its  melting  point.^  It  is  soluble  in 
liquids  containing  free  chlorine,  such  as  nitrate-chloride  mixtures 
of  acid  reaction,  and  to  a  lesser  degree  in  acid  ferric  chloride  solu- 
tions. These  facts  should  be  borne  in  mind,  in  order  to  avoid 
injury  to  platinum  vessels,  and  also  in  connection  with  the  pos- 
sible presence  of  platinum  in  certain  analytical  solutions. 

1  A  mixture  made  by  melting  together  equal  parts  of  vaseline  and  beeswax  is 
very  suitable. 

2  An  alloy  called  "palau,"  containing  80  per  cent  of  gold  and  20  per  cent  of 
palladium,  is  for  some  purposes  a  satisfactory  substitute  for  platinum.  It  melts, 
however,  at  1370°,  much  lower  than  platinum. 


INTRODUCTION  41 

Platinum  easily  alloys  with  most  metals,  and  for  that  reason  it 
should  not  ordinarily  be  heated  in  contact  with  metals,  or  with 
compounds  of  easily  reducible  metals,  —  never,  if  carbon  or 
reducing  gases  are  also  present.  When  heated  with  carbon, 
platinum  slowly  takes  this  up  and  becomes  brittle.  The 
crucible,  therefore,  should  not  be  heated  in  a  reducing  flame; 
the  flame  should  be  non-luminous,  and  so  adjusted  that  the 
crucible  is  wholly  above  the  inner  cone.  Burners  of  the  Meker 
type  are  very  suitable  for  the  heating  of  platinum  ware.  "  Un- 
known "  substances  should  not  be  heated  in  platinum  ware,  nor 
should  phosphates  or  arsenates  be  heated  under  reducing  condi- 
tions; these  elements,  as  well  as  phosphides  and  arsenides  are 
ruinous  in  their  action  upon  platinum. 

Platinum  ware  should  be  kept  clean  and  well  polished.  For 
this  purpose  moistened  precipitated  silica  is  very  satisfactory; 
it  removes  most  impurities  and  polishes  the  platinum  without 
the  loss  of  more  than  a  very  few  milligrams.  Potassium  bi- 
sulphate,  fused  within  the  vessel,  is  an  efficient  cleansing  agent ; 
but  the  melt  should  be  poured  out  while  still  liquid. 

Only  triangles  of  platinum,  quartz,  or  pipeclay  should  be  used 
with  platinum  crucibles,  and  hot  platinum  vessels  should  not 
be  set  down  on  foreign  metals.  The  crucibles  should  be  handled 
while  hot  only  with  platinum-shod  tongs.^ 

*  Modern  platinum  ware  is  often  inferior  in  quality  to  that  on  the  market  some 
years  ago,  and  the  cause  has  been  the  subject  of  special  inquiry  by  a  committee  of 
the  American  Chemical  Society.  The  main  objections  are :  "  (i)  Undue  loss  of 
weight  on  ignition ;  (2)  undue  loss  on  acid  treatment,  especially  after  strong  igni- 
tion ;  (3)  unsightly  appearance  of  the  surface  after  strong  ignition,  especially  after 
the  initial  stages  of  heating ;  (4)  adhesion  of  crucibles  and  dishes  to  triangles, 
sometimes  to  such  an  extent  as  to  leave  indentations  on  the  vessel  at  the  points 
of  contact  with  the  triangle,  even  when  complete  cooling  has  been  reached  before 
the  two  are  separated ;  (s)  alkalinity  of  the  surface  of  the  ware  after  strong  igni- 
tion; (6)  blistering;  and  (7)  development  of  cracks  after  continued  heating." 
It  is  the  general  opinion  that  the  trouble  arises  from  the  working  of  scrap  platinum 
into  chemical  ware.  The  main  difficulties  here  mentioned  are  not  characteristic  of 
platinum  ware  from  some  of  the  best  manufacturers. 

The  committee  recommends  that  purchasers  specify  that  platinum  ware  must 
show  no  marked  uneven  discoloration  on  heating,  must  give  no  test  for  iron  after 


42  QUANTITATIVE  CHEMICAL  ANALYSIS 

V.  THE  EVAPORATION  OF  LIQUIDS 

Evaporations  are  most  efficiently  carried  out  in  wide,  shallow 
dishes  or  casseroles,  which  expose  a  large  surface  of  liquid  to  the 
air.  Great  care  must  be  taken  to  prevent  the  loss  of  material  by 
spattering,  due  to  the  use  of  too  high  a  temperature,  or  from 
evaporation  to  a  low  volume  accompanied  by  the  crawling  of 
salts.  Solutions  evolving  gases,  or  which  are  to  be  boiled,  should 
always  be  covered.  And  liquids  which  contain  precipitates,  or 
other  solid  matter,  should  be  heated  cautiously  and  with  stir- 
ring ;  the  presence  of  solid  matter  at  the  bottom  of  a  liquid  often 
causes  violent  bumping,  and  this  may  even  result  in  the  de- 
struction of  the  vessel. 

The  evaporation  of  aqueous  solutions  rarely  requires  a  tem- 
perature higher  than  95-100° ;  the  use  of  the  steam  bath,  with 
which  mechanical  losses  due  to  boiling  and  bumping  need  not  be 
feared,  is  in  general  advisable  with  aqueous  solutions.  In  evap- 
orations on  the  steam  bath,  a  watch  glass  should  be  placed 
above  the  vessel  to  prevent  the  entrance  of  foreign  matter ;  but, 
by  means  of  a  glass  triangle,  it  should  be  elevated  above  the 
rim,  in  order  to  permit  the  free  exit  of  vapor. 

If  a  large  volume  of  liquid  is  to  be  evaporated,  the  dish  need 
not  necessarily  contain  it  all.  Fresh  portions  may  be  added 
from  time  to  time,  as  room  is  made  in  the  dish  by  evaporation. 

Liquids  should  be  transferred  from  one  vessel  to  another  only 
with  the  aid  of  a  glass  rod,  held  tightly  against  the  lip  of  the 
container.  In  the  case  of  vessels  without  lips,  in  order  to  pre- 
vent the  liquid  from  running  down  outside,  the  outside  edge 
should  be  coated  with  a  thin  film  of  vaseline. 

In  quantitative  work,  as  few  transfers  of  liquid  as  possible 

should  be  made.     Such  transfers  must  be  made  quantitative 

by  repeatedly  washing  the  former  container  with  small  portions 

of  wash  liquid ;  besides  being  the  most  effective,  this  method  of 

washing  also  keeps  the  total  volume  within  bounds. 

heating  for  two  hours,  and  that  the  rate  of  loss  per  hour  at  1100°  over  a  period 
of  four  hours  shall  not  exceed  0.2  mg. 


INTRODUCTION  43 

VI.  THE  VOLUMETRIC  MEASUREMENT  OF  LIQUIDS » 

In  a  well-equipped  weighing  room,  measurements  can  be  made 
with  a  precision  often  greater  than  is  necessary ;  for  the  errors  in- 
volved in  the  preparation  of  a  troublesome  precipitate  may  impair 
the  value  of  an  exact  weighing.  Although  the  measurement  of 
volume,  in  volumetric  analysis,  is  not  apt  to  be  so  precise  and 
rehable  as  the  measurement  of  weight,^  yet  the  results  of  volu- 
metric processes,  based  on  suitable  reactions,  are  frequently 
more  trustworthy  than  those  of  gravimetric  processes,  because 
the  volumetric  process  for  the  determination  of  the  substance 
is  less  liable  to  error.  With  proper  precautions  many  volu- 
metric processes  yield  excellent  results ;  and,  especially  in  techni- 
cal work,  where  time  is  an  essential  factor,  volumetric  processes 
are  very  often  used  in  preference  to  gravimetric.  In  order, 
however,  that  dangerous  errors  may  be  eliminated,  it  is  essential 
that  the  analyst  should  have  a  thorough  understanding  of  the 
precautions  necessary  for  the  attainment  of  a  high  degree  of 
accuracy. 

Volumetric  Apparatus.  The  exact  volumetric  measurement  of 
liquids  involves  the  use  of  certain  special  forms  of  apparatus. 

Burettes  are  graduated  glass  tubes  of  uniform,  small  diameter, 
for  the  measurement  of  variable  volumes  of  liquid  dehvered 
by  them  when  in  a  vertical  position.  The  outflow  is  controlled 
either  by  a  glass  stopcock,  or  by  means  of  a  short  rubber  tube 
provided  with  a  pinchcock  (or  other  suitable  device),  which 
connects  the  end  of  the  tube  with  a  glass  outlet,  or  tip.    The  glass 

*  For  more  detailed  information  on  this  subject,  see  Btdletin  of  the  Bureau  of 
Standards,  vol.  4,  pp.  553-601  (1908). 

^  Even  this  difficulty  can  be  obviated  by  the  use  of  weight  burettes;  i.e.  of  burettes 
of  such  construction  as  to  be  readily  weighable  both  before  and  after  the  removal 
of  the  quantity  of  solution  required  for  the  completion  of  the  given  reaction.  The 
difference  gives  the  weight  of  solution  required,  and,  provided  the  solution  has 
been  standardized  by  the  same  method,  the  quantity  of  the  substance  under  in- 
vestigation can  be  readily  calculated.  In  this  way,  if  based  on  suitable  reactions, 
exceedingly  exact  determinations  can  be  executed.  (A  similar  weight  burette 
should  always  be  used  as  a  counterpoise.) 


44  QUANTITATIVE  CHEMICAL  ANALYSIS 

stopcock  requires  the  use  of  a  lubricant,  as  vaseline,  to  permit 
easy  regulation;  and  the  rubber  joint  has  the  disadvantage 
that  it  is  attacked  by  some  solutions,  which  in  consequence 
sufifer  a  change  in  concentration,  —  e.g.  rubber  stopcocks  must 
be  avoided  with  permanganate  and  iodine  solutions. 

Transfer  Pipettes  are  very  narrow  tubes  which  are  designed 
to  deliver  specific  volumes  of  liquid;  an  enlargement  at  the 
center,  however,  greatly  reduces  the  length  required,  and  a 
mark  on  the  tube  above  this  enlargement  indicates  the  level  to 
which  the  instrument  must  be  filled  in  order  to  deliver  the  indi- 
cated volume  of  liquid.  Pipettes  are  filled  by  suction  and  are 
allowed  to  deliver,  from  a  vertical  position,  by  the  action  of 
gravity.  Owing  to  the  small  bore  of  the  tube,  the  pipette  is 
capable  of  a  high  degree  of  accuracy;  but  errors  of  manipula- 
tion may  render  the  measurements  inexact. 

Measuring  Flasks  are  usually  employed  in  the  measurement  of 
relatively  large  volumes.  The  neck  should  be  of  uniform  bore, 
should  extend  well  above  and  below  the  graduation  mark,  and 
should  be  small  enough  to  permit  of  accurate  reading,  but  suffi- 
ciently large  to  render  filling  and  emptying  easy.  For  the 
most  accurate  work,  the  flask  is  graduated  for  containing  the 
nominal  volume  of  liquid. 

Measuring  Cylinders  are  glass  vessels  provided  with  a  broad 
base,  or  foot,  and  with  a  lip  for  pouring.  They  are  graduated 
to  contain  or  deliver  variable  volumes  of  liquid,  and  are  used 
for  rough  measurements  only. 

Sources  of  Error  in  the  Use  of  Volumetric  Apparatus.  In 
the  use  of  volumetric  methods,  the  following  sources  of  error 
must  be  fully  reckoned  with  if  the  results  are  to  be  reliable. 

Water  in  the  Apparatus.  The  apparatus  must  nearly  al- 
ways be  washed  with  water  before  use,  and  any  water  retained 
would  of  course  alter  the  concentration  of  the  solution  to  be 
measured.  While  the  apparatus  might  be  dried  before  use,  it 
is  far  more  convenient,  and  just  as  accurate,  to  wash  it  out  with 
small  portions  of  the  solution,  and  to  discard  the  washings. 


INTRODUCTION  45 

Drainage  or  Afterflow.  When  an  aqueous  liquid  is  permitted 
to  flow  rapidly  from  a  burette,  or  pipette,  some  of  the  liquid 
adheres  to  the  inner  surface,  and  only  gradually  runs  down  to 
the  liquid  below.  To  avoid  errors  from  this  source,  the  rate  of 
outflow  must  be  limited  by  the  size  of  the  outlet,  or  a  sufficient 
time  interval  must  be  allowed  to  elapse  before  the  reading  is 
made. 

In  the  case  of  transfer  pipettes,  the  outlets  should  be  made 
of  such  size  that  the  free  outflow  shall  require  at  most  i  minute, 
but  not  less  than  15,  20,  and  30  seconds,  respectively,  for  5,  10, 
and  50  pipettes.  The  free  outflow  from  a  burette  should  never 
take  much  more  than  3  minutes,  nor  less  than  90  and  50  seconds 
for  50  and  30  cc.  burettes,  respectively.  Burette  and  pipette 
tips  should  be  made  with  a  gradual  taper  of  2-3  cm. ;  a  sudden 
contraction  at  the  orifice  is  not  permissible,  and  the  tip  should 
be  well  finished. 

Grease  Films.  Certain  substances,  especially  grease,  adher- 
ing to  the  walls  of  measuring  vessels  prevent  their  uniform 
wetting ;  the  tendency  of  water  to  collect  into  droplets,  instead 
of  flowing  uniformly  over  the  glass,  indicates  the  presence  of  a 
film  of  grease.  Imperfect  wetting  not  only  distorts  the  meniscus, 
but,  by  influencing  the  residue  adherent  to  the  walls,  it  also  gives 
rise  to  irregularities  in  the  delivery  of  liquid. 

When  cleaning  accurately  calibrated  measuring  vessels  with 
sulphuric  acid-dichromate  solution,  they  should  be  filled  with 
the  cold  solution  and  allowed  to  stand  overnight,  or  longer. 
The  use  of  hot  liquids  in  such  vessels  is  to  be  avoided  on  account 
of  the  possible  thermal  after  effect  on  the  glass ;  while  for  cer- 
tain kinds  of  glass  this  effect  may  be  negligible,  we  have  no  as- 
surance that  it  is  so. 

Parallax.  In  apparatus  in  which  the  volume  is  limited  by  a 
meniscus,  the  reading  or  setting  is  made,  when  possible,  on  the 
lowest  point  of  the  meniscus.  This  point  lies  at  the  center 
of  the  tube;  the  reading  should  be  made,  therefore,  with 
the  tube  in  a  vertical  position,  and  with  the  eye  so  located 


46 


QUANTITATIVE  CHEMICAL  ANALYSIS 


that  the  line  of  sight  is  perpendicular  to  the  main  axis  of  the 
tube. 

In  making  this  reading,  it  is  well  to  place  a  dark-colored  screen 
immediately  below  the  meniscus;  this  will  render  the  profile 
of  the  meniscus  dark  and  clearly  visible  against  a  light  back- 
ground. A  convenient  device  for  this  purpose  is  a  collar-shaped 
section  of  black  rubber  tubing,  cut  open  at  one  side,  and  of  such 
a  size  as  to  clasp  the  tube  firmly. 

Variations  in  Temperature.  The  volume  occupied  by  a  given 
weight  of  water,  as  well  as  the  capacity  of  the  measuring  vessel, 
is  dependent  upon  the  temperature ;  and  the  error  involved  in 
the  measurement  of  the  volume  of  a  given  mass  of  water,  at  any 
other  temperature  than  the  standard  one,  is  due  to  the  joint 
effect  of  the  changed  capacity  of  the  vessel  and  the  changed  vol-  ^ 
ume  of  the  liquid. 

The  coefficient  of  cubical  expansion  of  ordinary  glass  is  very 
close  to  0.000025  ;  but  the  volume  change  of  the  water  is  much 
greater  than  that  of  the  glass  measuring  vessel,  and  also  much 
less  uniform  from  degree  to  degree.  The  factors  by  which  a 
volume  of  water,  measured  at  temperatures  ranging  from  10- 
29°  in  a  vessel  calibrated  for  20°,  must  be  multipHed  in  order  to 
obtain  the  true  volume  occupied  by  the  liquid  at  20°,  are  given 
in  the  following  table : 


Temperature  of  the  Water 

\  Units 
Tens\ 

0 

1 

2 

3 

4 

1 

1. 001 24 

I.OOII7 

1. 00109 

1. 001 00 

1.00089 

2 

I.OOOOO 

0.99981 

0.99961 

0.99941 

0.99919 

Temperature  op  the  Water                                   j 

\  Units 
Tens  \^ 

6 

6 

7 

8 

9 

1 

1.00077 

1.00064 

1.00049 

1.00034 

1. 000 1 8 

2 

0.99896 

0.99873 

0.99848 

0.99822 

0.99798 

INTRODUCTION  47 

If  the  prevailing  temperature  does  not  differ  by  more  than, 
say,  3°  from  the  standard,  this  correction  may  ordinarily  be 
omitted.  In  the  case  of  solutions  of  0.2  N  concentration,  or  less, 
the  corrections  differ  so  little  from  those  for  pure  water  that  the 
factors  given  in  the  table  may  be  used  without  appreciable  error. 

In  order  to  illustrate  the  use  of  such  factors,  let  us  suppose 
that,  in  the  standardization  of  a  solution,  a  burette  graduated 
correctly  for  20°  is  used  at  an  actual  temperaure  of  27°,  and 
that  the  indicated  volume  of  solution  withdrawn  is  28.75  cc. 
Then  the  true  volume  at  20°  of  this  quantity  of  liquid  is  28.75  X 
0.99848=28.70  cc.  And,  if  a  determination  is  later  made  with 
this  solution  at,  say,  17°,  and  the  indicated  volume  used  is 
28.68  cc,  then  the  true  volume  at  20°  is  28.68  X  1.00049=  28.70  cc. 
That  is  to  say,  the  same  quantity  of  reagent  is  contained  in  an 
apparent  volume  of  28.75  cc.  at  27°,  or  in  an  apparent  volume 
of  28.68  cc.  at  17°,  as  is  contained  in  an  actual  volume  of  28.70  cc. 
of  the  solution  at  20°.^ 

Different  Units  of  Volume.  Unfortunately,  a  number  of 
different  "  liters  "  have  been  suggested  for  use  in  volumetric 
analysis.  The  normal  liter,  that  is,  the  volume  occupied  by  a 
kilogram  of  water,  weighed  in  a  vacuum  and  measured  at  4°, 
would  manifestly  be  out  of  the  question  if  it  had  to  be  deter- 
mined in  that  way.  The  so-called  "  Mohr  liter  "  is  the  volume 
occupied  by  a  kilogram  of  water  when  weighed  in  the  air  with 
brass  weights  at  a  temperature  of  17.5° ;  but  this  volume  varies 
with  the  atmospheric  conditions.  Other  "  Hters "  involving 
measurements  at  15°,  15.5°,  or  20°  have  been  suggested  by 
various  chemists. 

It  matters  Httle,  in  analytical  work,  which  liter  is  adopted, 
but  it  is  of  the  greatest  importance  to  have  the  pipettes,  burettes, 
and  measuring  flasks  rigorously  consistent  with  one  another.  This 
matter  requires  especial  emphasis,  since  apparatus,  if  not  speci- 

^  At  27**  the  actual  volume  of  the  liquid  measured  is  a  shade  greater  than  28.75  cc, 
and  at  17°  it  is  a  shade  less  than  28.68  cc.  The  measuring  vessel  is  larger  at  27°, 
and  smaller  at  17°,  than  it  is  at  20°, 


48  QUANTITATIVE  CHEMICAL  ANALYSIS 

fically  ordered,  may  be  supplied  by  dealers,  at  different  times, 
graduated  according  to  different  systems;  and  mixed  gradua- 
tions may  thus  come  into  the  hands  of  an  individual  analyst. 
As  an  example  of  the  magnitude  of  the  errors  which  might  thus  be 
introduced,  it  should  be  noted  that  the  normal  liter  is  related 
to  the  Mohr  liter  as  looo :  1002.3. 

Much  of  the  graduated  apparatus  on  the  market  bears  no 
mark  by  means  of  which  the  unit  of  volume  represented  can  be 
recognized,  and  even  when  this  is  clearly  designated  the  per- 
centage error  represented  may  be  large.  It  is  not  advisable, 
therefore,  to  use  any  piece  of  graduated  apparatus,  unless  its 
actual  value  is  well  known. 

Owing  to  the  great  difiiculty  in  measuring  directly  the  re- 
lation between  cubic  capacity  and  the  unit  of  length,  the  Inter- 
national Committee  of  Weights  and  Measures  defines  the  liter 
as  "  the  volume  occupied  by  the  mass  of  one  kilogram  of  pure 
water  at  its  maximum  density  under  normal  atmospheric  pres- 
sure." This  is  almost  exactly  1000  cc.^  and  for  all  practical 
purposes  may  be  regarded  as  such. 

It  is  now  customary  to  use  this  true  liter  as  the  standard, 
but  of  course  it  is  out  of  the  question  to  weigh  a  kilogram  of 
water  at  4°  in  a  vacuum ;  some  convenient  temperature  — 
preferably  the  average  working  temperature  of  the  laboratory 
—  must  be  selected,  and  the  necessary  corrections  made.  If  a 
liter  flask  is  marked  correctly  at  20°,  this  means  that  at  20°  it 
will  contain  a  mass  of  water  (998.234  g.)  which  occupies  a  volume 
equal  to  that  occupied  by  1000  grams  of  pure  water  at  4°.  This 
quantity  of  water,  if  weighed  with  brass  weights  in  air  of  mean 
humidity,  at  20°  and  760  mm.,  has  an  apparent  weight  of  997.18 
grams. 

The  Calibration  of  Volumetric  Apparatus.  The  weight  of 
brass  (brass  weights)  which  will  be  required  to  counterbalance 
one  liter  of  pure  water  must  be  calculated  from  the  temperature 
of  the  water  and  the  density  of  the  air.     The  following  table 

*  About  1000.029  cc, 


INTRODUCTION 


49 


indicates,  for  temperatures  of  the  water  (and  room)  ranging  from 
15-29°,  how  many  milligrams  less  than  1000  grams  a  quantity  of 
water  will  weigh  which  is  sufficient  to  fill  to  the  mark  a  i -liter  flask 
correctly  calibrated  for  20°,  the  weighing  being  carried  out  in  air 
of  50  per  cent  humidity  at  760  mm.  pressure  (unreduced).^ 


\  Units 
Tens\^ 

0 

1 

2 

3 

4 

5 

6 

7 

8 

9 

1 

— 

~^ 

— 

— 

— 

1950 

2100 

2260 

2440 

2620 

2 

2820 

3030 

3240 

3470 

3710 

3960 

4210 

4480 

4760 

5040 

If  such  a  flask  is  filled  to  the  mark  with  water  of  22.4°,  for 
example,  the  water  will  under  the  conditions  of  the  table  require 
a  counterpoise  of  1000-3.332  =  996.668  grams. 

The  determination  of  the  capacity  of  a  measuring  flask  is 
carried  out  by  weighing  the  water  contained  in  it,  while  the 
volume  of  water  delivered  by  a  burette  or  pipette  is  determined 
by  weighing  this  water  after  its  delivery  into  another  vessel. 
The  temperature  of  the  water,  which  should  be  the  same  as  that 
of  the  room,  should  be  taken  immediately  before  and  after  the 
experiment. 

In  the  calibration  of  a  flask,  the  dry  flask  is  placed  upon  the 
right-hand  pan  of  the  balance,  together  with  the  nominal  weight 
of  its  capacity,  i.e.  with  as  many  grams  as  it  is  supposed  to  con- 
tain cubic  centimeters,  and  then  tare  material  is  placed  upon 
the  left-hand  pan  until  the  balance  is  brought  into  equilibrium. 
The  weights  are  then  removed  from  the  right-hand  pan,  the 
flask  is  filled  to  the  mark  with  water,  and  weights  are  added 
until  the  balance  is  again  in  equilibrium.  The  nominal  capacity 
weight,  minus  the  additional  weight  which  is  required  upon  the 
right-hand  pan  in  order  to  reestablish  equilibrium,  is  equal  to 
the  weight  of  the  water  in  the  flask. 

In  the  case  of  burettes  and  pipettes,  a  covered  beaker  is  placed 

1  These  values  depend  upon  the  specific  gravities  of  the  water  and  the  brass, 
the  density  of  the  air,  and  the  coefficient  of  cubical  expansion  of  the  glass 
vessel. 


50  QUANTITATIVE  CHEMICAL  ANALYSIS 

upon  the  right-hand  pan,  together  with  the  nominal  weight  in 
grams  of  the  volume  to  be  deHvered,  after  which  the  balance 
is  brought  into  equilibrium  by  the  addition  of  tare  material 
to  the  left-hand  pan.  The  water  is  then  allowed  to  run  into  the 
beaker,  which  is  replaced  upon  the  right-hand  pan.  The  subse- 
quent procedure  is  the  same  as  that  described  above. 

The  difference  between  the  additional  weight  required  to 
reestablish  equilibrium  and  that  calculated  from  the  above 
table  indicates  directly  the  error  of  the  vessel.  If,  for  example, 
it  be  found  necessary,  in  order  to  reestabhsh  equilibrium  in  the 
case  of  a  500-cc.  flask  filled  to  the  mark  with  water  of  22.4°, 
to  add  1.832  g.,  instead  of  the  calculated  1.666  g.,  then  it  follows 
that  the  vessel  is  0.166  cc.  too  small.  On  the  other  hand,  in  test- 
ing the  25  cc.  segment  of  a  30-cc.  burette,  at  a  temperature  of  17°, 
the  additional  weight  required  on  the  right-hand  pan  should  be 
(25X226o)-Mooo=57  mg. ;  if,  instead  of  this,  it  be  found  that 
an  additional  weight  of  15  mg.  is  required  on  the  left-hand  pan, 
then  the  25  cc.  segment  is  57  mg.  — (  — 15  mg.)  =  0.072  cc.  too 
large.^ 

In  the  caHbration  and  use  of  burettes,  the  liquid  should  in 
general  be  allowed  to  flow  from  the  zero  mark  to  some  second 
level  in  the  burette. 

D.  THE  PREPARATION  OF  SAMPLES  FOR  ANALYSIS 

It  is  not  easy  to  give  general  rules  for  the  preparation  of  sub- 
stances for  analysis,  because  it  is  necessary  to  proceed  differ- 
ently in  different  cases.  In  all  cases,  however,  the  samples 
should  promptly  be  transferred  to  tightly  stoppered  bottles  or 
weighing  tubes. 

In  technical  analyses,  for  the  purpose  of  determining  the 
commercial  value  of  an  article,  or  of  controlling  processes  of 
manufacture,  materials  must  be  analyzed  as  they  are.  But, 
in  every  case,  especial  care  should  be  taken  to  make  up  a  sample 

*  This  error  is  too  large  to  be  tolerated. 


INTRODUCTION  51 

which  will  represent  as  nearly  as  possible  the  average  composition 
of  the  whole  lot. 

If,  on  the  other  hand,  it  is  desired  to  determine  the  atomic 
composition  of  a  compound,  it  is  necessary  to  select  or  prepare 
pure  material  for  analysis.  This  may  seem  simpler  than  it 
really  is.  Many  compounds  absorb  or  give  up  moisture  upon 
exposure  to  the  air,  and  their  treatment  should  vary  with  their 
nature,  as  illustrated  in  the  following  cases.  Salts  such  as 
Na2S04 .  10  H2O  and  Na2C03 .  10  H2O,  which  effloresce  in 
ordinary  air,  may  be  dried,  after  recrystallization,  by  strongly 
pressing  the  powdered  material  between  several  layers  of  filter 
paper,  the  paper  being  renewed  as  long  as  moisture  continues 
to  be  taken  up ;  MgS04 .  7  H2O  and  NaKC4H406 .  4  H2O, 
which  do  not  lose  water  of  constitution  in  ordinary  air,  may  be 
spread  out  upon  filter  paper,  covered  with  another  sheet,  and 
allowed  to  dry  at  the  ordinary  temperature.  Compounds  such 
as  HFe(S04)2  •  4  H2O  and  CaC4H406 .  H2O,  which  do  not 
effloresce  in  artificially  dried  air,  but  which  undergo  chemical 
change  at  icx5°,  may  be  conveniently  dried  in  a  desiccator,  over 
calcium  chloride.  Substances,  as  KHC4H4O6,  sugar,  etc.,  which 
readily  give  up  hygroscopic  moisture  at  100°,  without  other 
alteration,  are  best  dried  in  an  oven  at  that  temperature ;  while 
K2PtCl6,  which  retains  moisture,  or  dries  only  slowly  at  100°, 
but  which  decomposes  below  a  red  heat,  should  be  dried  in  an 
oven  at,  say,  130°.  Finally,  substances  such  as  NaCl,  Na2S04, 
etc.  may  be  given  a  preliminary  drying,  in  a  covered  vessel,  at 
130°,  or  higher,  to  prevent  decrepitation,  and  then  be  ignited, 
more  or  less  strongly,  depending  upon  their  nature.  In  every 
case,  the  sample  should  be  dried,  without  decomposition,  to 
constant  weight. 

Substances  used  in  testing  the  accuracy  of  analytical  pro- 
cesses, or  in  standardizing  volumetric  solutions,  must  also  be 
extremely  pure.  In  fact,  compounds  are  generally  favored 
which  are  non-hygroscopic,  and  which  may  readily  be  prepared 
in  a  pure  condition ;   if  possible,  it  is  well  to  select  compounds 


52  QUANTITATIVE  CHEMICAL  ANALYSIS 

which  normally  do  not  contain  water  of  crystallization.  Many 
salts  can  be  obtained  sufficiently  pure  in  the  market ;  but  their 
purity  should  never  be  accepted  on  faith.  If  tests  indicate  the 
presence  of  impurity,  and,  often,  if  the  salt  contains  water  of 
crystallization,  the  material  should  be  recrystallized. 

For  this  purpose,  a  convenient  weight  of  the  salt  is  dissolved 
in  the  least  possible  quantity  of  hot  water,  using  a  quantity  of 
water  not  quite  sufficient  to  dissolve  the  whole  lot;  the  hot 
solution  is  poured  into  a  fluted  filter,  held  in  a  stemless  funnel, 
and  the  filtrate  is  received  with  continuous  stirring  in  a  beaker, 
which  itself  is  immersed  in  cold  water,  in  a  larger  vessel.  The 
rapid  cooling  and  constant  stirring  cause  the  formation  of  a 
fine  crystaUine  powder,  which  is  almost  free  from  inclosed 
mother-liquor.  The  crystalline  powder  is  filtered  off  in  a  funnel 
containing  a  perforated  platinum  cone,  the  adhering  mother- 
liquor  being  removed  by  suction  or  in  a  centrifuge.  Two  such 
recrystallizations  will  nearly  always  suffice.  According  to  the 
nature  of  the  substance,  it  is  dried  in  the  air  at  a  specific  tem- 
perature, or  in  a  desiccator,  to  constant  weight. 

Concerning  the  preparation  of  samples  for  analysis  by  beginners 
in  quantitative  analysis,  the  reader  should  consult  the  Appendix. 


PART   II 

GRAVIMETRIC  ANALYSIS 
EXERCISES  WITH  THE  BALANCE 

Before  beginning  work  at  the  balance,  read  carefully  the 
rules  given  on  pp.  9-1 1  of  Part  I,  and  observe  them  always. 

Determination  of  the  Rest-Point.  Determine  the  rest-point 
of  the  unloaded  balance,  according  to  the  method  on  p.  11.  If 
this  is  not  more  than  one  division  from  the  center  of  the  scale, 
the  balance  may  be  used  by  the  student;  otherwise  it  will  be 
adjusted  by  an  instructor,  upon  request.  The  beginner  should 
not  attempt  this  adjustment. 

Determination  of  the  Weight  of  an  Object.  Clean  two  porce- 
lain crucibles,  rinse  them  with  distilled  water,  and  allow  them 
to  drain.  Place  each  crucible  upon  a  pipestem  triangle,  sup- 
ported upon  a  tripod,  and  heat  with  the  colorless  flame  of  a 
Bunsen  or  Tyrill  burner,  —  gently  at  first,  and  then  to  a  red 
heat.  Allow  the  crucibles  to  cool  off  somewhat,  but  while  still 
warm,  place  them  in  a  desiccator,  using  the  crucible  tongs. 
(A  piece  of  apparatus  which  has  been  ignited,  or  dried,  for  weigh- 
ing must  not  be  touched  by  the  hand  before  it  has  been  weighed.)  Al- 
low the  crucibles  to  cool  in  the  desiccator  for  at  least  20  minutes. 

Now  with  the  crucible  tongs  place  a  crucible  on  the  left-hand 
pan  of  the  balance,  and,  by  means  of  the  forceps  in  the  weight 
box,  place  weights  on  the  right-hand  pan  until  they  balance  the 
crucible  to  within  0.005  g.  Begin  with  a  weight  which  you 
think  will  approximately  balance  the  object,  lower  the  balance 
beam,  and  gently  release  the  pan  supports.  It  will  then  be  seen 
which  side  is  the  heavier.     Finally  adjust  the  rider,  so  that, 

S3 


54  QUANTITATIVE  CHEMICAL  ANALYSIS 

when  the  beam  is  swinging  freely,  the  pointer  swings  equal 
distances  on  both  sides  of  the  rest-point.  Always  try  the 
weights  in  the  order  in  which  they  occur  in  the  box,  beginning 
with  the  heavier  ones,  and  using  the  rider  for  weights  smaller 
than  5  or  lo  milHgrams,  according  to  the  number  of  large  divi- 
sions on  the  beam. 

As  soon  as  the  object  appears  to  be  balanced,  raise  and  lower 
the  beam,  and  make  another  observation.  Read  the  weight  of 
the  crucible  hy  noting  in  order  the  vacant  spaces  in  the  box,  begin- 
ning with  the  largest  missing  weight;  and  check  this  reading  as  the 
weights  are  returned  to  the  box.  Be  sure  also  to  note  the  weight 
recorded  by  the  rider,  and  then  lift  it  from  the  beam.  Always 
record  the  weight,  in  pencil  as  it  is  first  read,  and  in  ink  after  it 
has  been  checked,  and  always  in  the  record  book. 

In  this  manner,  weigh  the  two  crucibles  separately,  and  then 
weigh  them  together,  entering  all  three  results  in  the  notebook. 
{In  connection  with  the  keeping  of  records,  see  the  remarks  on 
p.  5.)  The  sum  of  the  separate  weights  should  agree  closely 
with  the  result  obtained  upon  weighing  both  crucibles  together, 
—  within,  say,  0.0002  g. 


GRAVIMETRIC  ANALYSIS  55 

THE  DETERMINATION  OF  CHLORINE  IN  A  SOLUBLE 
CHLORIDE 

The  sample  may  be  pure  sodium  chloride,  or  it  may  be  an  artifi- 
cially prepared  mixture  of  sodium  chloride  and  sodium  carbonate. 

Method.  The  aqueous  solution  of  the  chloride  is  acidified 
with  nitric  acid  and  treated  with  silver  nitrate  in  excess.  The 
chlorine  is  quantitatively  precipitated  as  silver  chloride,  which 
is  filtered  off,  washed,  dried,  and  weighed.  Other  acids  which 
yield  silver  salts  insoluble  in  nitric  acid  must  of  course  be  absent. 

A.  A  Paper  Filter  is  Used:  Procedure.  Carefully  wipe  off 
the  stoppered  sample  tube,  without  directly  touching  it,  and 
weigh  it  accurately  to  o.i  mg.,  entering  the  weight  in  the  note- 
book (see  p.  5).  With  the  tube  held  just  above  a  clean  300  cc. 
beaker  (labelled  "  I  "),  remove  the  stopper,  and,  without  loss  of 
material  in  the  form  of  dust  or  otherwise,  carefully  transfer  to 
the  beaker  0.2—0.3  g.  of  the  sample.  Replace  the  stopper, 
weigh  again  to  o.i  mg.,  and  enter  the  result  in  the  notebook. 
The  first  weight  minus  the  second  gives  the  weight  of  sample 
taken.  Weigh  out  a  second  0.2—0.3  g.  portion  into  another 
beaker  (labelled  "  II  ")>  entering  the  data  as  before.  (Be  sure 
always  to  cover  the  beakers  promptly.) 

Treat  each  sample  as  follows:  Add  100-125  cc.  of  distilled 
water,  stir,  and  acidify  with  nitric  acid;  add  the  acid  slowly 
with  stirring  until  blue  litmus  paper  indicates  an  acid  reaction 
when  touched  with  the  moist  stirring  rod.  Assuming  the  ma- 
terial to  be  sodium  chloride,  add  slowly  with  stirring  a  volume 
of  silver  nitrate  solution  4-5  cc.  in  excess  of  the  calculated  quan- 
tity (fiP&te  strength  of  the  reagents,  see  the  Appendix).  Cover 
the  beaker  with  a  watch  glass,  and,  with  occasional  stirring, 
slowly  heat  the  solution  to  boiling.  This  treatment  will  cause  the 
ipecipitate  to  coagulate  and  settle,  leaving  the  liquid  clear.  The 
precipitate  should  not  be  exposed  to  direct  sunlight,  and  great 
care  should  be  taken  to  avoid  the  loss  of  material.  Finally, 
add  a  little  silver  nitrate  solution  to  the  hot  supernatant  liquid, 


56  QUANTITATIVE  <:HEMICAL  ANALYSIS 

to  test  for  complete  precipitation;   if  a  precipitate  forms,  add 
5  cc.  more,  stir,  allow  to  settle,  and  test  again. 

Prepare  two  9  cm.  "  ashless  "  filters  (see  p.  31),  and  decant 
the  hot  liquid  through  the  filter  in  each  case,  leaving  the  pre- 
cipitate mostly  in  the  beaker.  The  filtrate  should  be  free  from 
turbidity ;  otherwise  refilter  it  through  the  same  paper.  (Fi- 
nally, pour  the  clear  filtrate  into  the  bottle  for  "  Silver  Residues.'') 
With  clean  receptacles  for  the  washings,  wash  each  precipitate 
twice  by  decantation  with  10  cc.  portions  of  hot  water  (to  which 
it  is  well  to  add  3-4  drops  of  nitric  acid),  running  the  washings 
through  the  filters,  and  then  transfer  each  precipitate  to  the 
corresponding  filter.  This  transfer  is  made  by  means  of  a  stream 
of  hot  water  from  the  wash  bottle,  the  adhering  particles  being 
loosened  with  the  aid  of  a  "  policeman  "  (see  p.  33).  Wash  the 
precipitates  and  filters  with  hot  water  until  5  cc.  of  the  washings 
show  no  turbidity  whatever  with  i  drop  of  6-normal  hydrochloric 
acid.  After  the  complete  drainage  of  the  filters,  cover  each 
funnel  with  an  ordinary  filter  paper,  crimping  its  edges  over  the 
rim  of  the  funnel,  and  place  it  (numbered  according  to  the  note- 
book record,  and  bearing  the  student 'j^ame  and  desk  number) 
in  a  drying  oven,  at  90-100° ;  allow  the  filters  and  precipitates 
to  dry  completely. 

Now,  in  a  quiet  place,  open  each  filter  over  a  smoothly  trimmed 
square  of  black  glazed  paper,  about  6  inches  in^iameter,  and 
with  great  care  to  avoid  loss  (by  finally  rubbing  tmifolds  of  the 
paper  gently  together),  transfer  the  precipitate  tcAthe  center 
of  the  glazed  paper ;  do  not  rub  off  any  fibers  from  tfte  paper. 
Cover  the  material  on  the  paper  with  an  inverted  funnel  or 
watch  glass  for  protection.  'li(^^^ 

With  the  paper  refolded  flat,  bend  the  top  over  and  aH^the 
fold ;  then  roll  the  paper  into  a  small  bundle  and  place  i^^^ 
weighed  porcelain  crucible.  Place  the  crucible  on  its  side,  up^ 
a  triangle,  and  ignite  gently  until  the  paper  is  decomposed. 
Then,  with  the  flame  at  the  back  of  the  crucible,  ignite  strongly 
until  the  carbon  is  consumed.     (See  Part  I,  pp.  38-39.)    After 


GRAVIMETRIC  ANALYSIS  57 

cooling,  add  2  drops  of  6-nonnal  nitric,  and  i  of  hydrochloric 
acid,  and  heat  gently,  with  great  caution  to  avoid  loss,  until  the 
acids  are  expelled.     Allow  to  cool. 

Place  the  cold  crucible  upon  a  piece  of  glazed  paper  (to  catch 
any  material  that  might  be  spilled),  and  transfer  to  it  the  main 
quantity  of  the  silver  chloride,  brushing  the  last  traces  of  the 
salt  into  the  crucible  with  a  small  camel's  hair  brush.  Now  add 
2  drops  of  nitric,  and  i  of  hydrochloric  acid,  expel  the  acids 
(observe  caution),  and  then  slowly  raise  the  temperature  until 
the  salt  just  begins  to  fuse ;  at  that  point,  remove  the  flame. 
Place  the  warm  crucible  in  a  desiccator,  allow  it  to  cool,  and  then 
weigh  it.  Without  the  acids,  repeat  the  heating,  etc.  until  the 
last  change  in  weight  observed  does  not  exceed  0.2  mg.  (Fi- 
nally, place  the  silver  salt  in  the  bottle  for  silver  residues.) 

From  the  weight  of  the  silver  chloride  yielded  by  each  sample, 
calculate  the  percentage  content  of  chlorine. 

Notes.  —  i.  The  solution  is  acidified  with  nitric  acid  in  order  to  prevent 
the  formation  and  precipitation  of  substances,  as  silver  oxide,  carbonate, 
phosphate,  etc.,  which  are  insoluble  in  water  but  soluble  in  nitric  acid. 
The  acid  also  helps  to  coagulate  the  precipitate. 

2.  It  is  safer  not  to  boil  the  acidified  solution  until  the  silver  nitrate 
has  been  added ;  otherwise  a  trace  of  hydrochloric  acid  might  be  oxidized 
by  nitric  acid  and  the  liberated  chlorine  partially  volatilized.  While  not  at 
all  likely  at  so  low  an  acid  concentration,  it  is  nevertheless  needless  to  run 
the  risk  of  this  possible  action. 

3.  Silver  chloride  is  decomposed  sHghtly  by  direct  sunlight,  with  the 
loss  of  chlorine ;  this  action  is  much  less  in  diffused  light,  and  can  be  remedied 
by  the  acid  treatment  indicated.  In  diffused  daylight,  however,  the  error 
from  this  source  is  really  too  insignificant  to  affect  the  accuracy  of  the 
determination. 

4.  The  precipitate  is  washed,  before  drying,  in  order  to  free  it  from  the 
non-volatile  soluble  salts  present  in  the  solution.  The  washing  is  to  be 
discontinued  as  soon  as  5  cc.  of  the  washings  give  no  turbidity  upon  the 
addition  of  chloride-ion ;  only  one  or  two  drops  of  hydrochloric  acid  should 
be  used  in  this  test,  owing  to  the  solubility  of  silver  chloride  in  concentrated 
chloride  solutions.  The  wash  water  should  be  hot  in  order  to  keep  the 
precipitate  coagulated ;  this  is  still  more  assured  if  the  water  is  very  slightly 
acidified  with  nitric  acid. 


58  QUANTITATIVE  CHEMICAL  ANALYSIS 

5.  Silver  chloride  itself  is  practically  insoluble  in  cold  water;  slightly 
over  I  mg.  per  liter  {i.e.  i  part  per  million  of  water)  at  20°.  Owing  to  the 
common-ion  effect,  the  solubility  is  still  less  in  the  presence  of  a  slight 
excess  of  either  silver-,  or  chloride-ion,  —  e.g.  in  very  dilute  silver  nitrate 
or  hydrogen  chloride  solution.  In  hot  water  the  salt  is  more  soluble  (about 
22  parts  per  million,  at  100°) ;  but,  fortunately,  the  speed  of  solution  is  so 
slow  that  the  precipitate  may  be  thoroughly  washed  with  hot  water,  without 
undue  error. 

6.  Silver  chloride  should  not  be  ignited  with  the  paper,  because  it  would 
be  largely  reduced  to  metal  (AgCl+"H,"  from  the  paper=Ag+HCl). 
The  small  quantity  which  must  necessarily  be  ignited  with  the  paper  is 
probably  all  reduced ;  but  it  can  with  certainty  be  changed  back  to  chloride, 
without  loss. 

Silver  chloride  is  also  reduced  by  zinc  and  dilute  sulphuric  acid  to  metallic 
silver,  and  the  salt  adhering  to  the  crucible  after  ignition  can  be  removed 
by  means  of  this  action.  Upon  fusion  with  sodium  carbonate,  owing  to  the 
unstable  nature  of  silver  carbonate  and  oxide,  silver  chloride  readily  yields 
metallic  silver. 

7.  Like  most  substances,  silver  chloride  is  appreciably  volatile  somewhat 
above  its  melting  point.  It  is  not  safe,  therefore,  to  fuse  the  precipitate 
over  an  ordinary  burner;  there  is  too  much  danger  of  overheating.  If 
heated  slowly,  however,  only  to  incipient  fusion  (around  the  edges),  no 
error  need  be  feared  at  this  point. 

8.  Silver-ion  has  a  great  tendency  to  enter  into  the  formation  of  complex 
ions,  as  [Ag(NH3)2]"^,  [Ag(CN)2]",  [AgSaOa]",  and,  in  consequence,  the  so- 
called  insoluble  salts  of  silver  often  enter  into  solution  upon  the  addition 
of  ammonia,  or  of  an  alkali  cyanide  or  thiosulphate.  Such  facts  as  these 
must  be  acquired  and  kept  in  mind  by  the  analytical  chemist ;  chemical 
knowledge  can  be  of  the  greatest  service  in  combating  difficulties,  while 
lack  of  knowledge  is  often  the  occasion  of  trouble. 

9.  The  silver  residues  should  conscientiously  be  turned  in  by  the  student. 
Assuming  that  dupHcate  determinations  require  on  the  average  55  cc.  of 
0.2-normal  silver  nitrate  solution,  and  taking  into  account  the  analyses 
which  have  to  be  repeated,  the  residues  returnable  by  a  class  of  100  students 
should  contain  about  150  g.  (5  oz.)  of  metallic  silver.  This  can  easily  be 
recovered  and  transformed  into  pure  silver  nitrate,  for  future  use  by  other 
students. 

B.  A  Gooch  Crucible  is  Used :  Procedure.  Weigh  out  two 
samples  of  the  substance,  of  about  0.25  g.  each,  and  convert  the 
chloride  into  silver  chloride  as  described  in  procedure  A.     Mean- 


GRAVIMETRIC  ANALYSIS  59 

while,  prepare  two  Gooch  crucibles,  following  the  directions  on 
pp.  33-35 ;  and  finish  the  analysis  according  to  the  details  there 
given. 

Notes.  —  i.  Silver  itself  (e.g.  as  AgCl),  as  well  as  bromides,  iodides, 
cyanides,  thiocyanates,  etc.,  may  be  determined  in  a  similar  manner,  with 
the  use  of  Gooch  crucibles.  The  disadvantage  of  using  a  paper  filter,  e.g, 
with  silver  bromide  or  cyanide  (requiring  the  conversion  of  reduced  silver 
back  to  the  bromide  or  cyanide),  is  at  once  apparent. 

2.  Chlorates,  etc.,  may  be  determined  as  silver  halides  by  first  reducing 
them  to  the  oxygen-free  salts,  and  precipitating  these  with  silver  nitrate. 

3.  For  the  principles  involved  in  the  determination  of  these  substances 
when  two  or  more  of  them  are  present  in  the  same  sample,  the  student  is 
referred  to  Part  V,  Problems  37,  38,  90,  91,  and  93. 


6o  QUANTITATIVE  CHEMICAL  ANALYSIS 

THE    DETERMINATION   OF   IRON   AND   OF   SULPHUR   IN   A 
SOLUBLE  SULPHATE  OF  IRON 

The  sample  may  be  pure  ferrous  ammonium  sulphate,  pure 
ferric  alum,  or  an  artificially  prepared  mixture  of  anhydrous 
ferric  sulphate,  sodium  carbonate,  and  potassium  sulphate. 
This  mixture  is  readily  soluble  in  dilute  hydrochloric  acid. 

Method.  The  sample  is  dissolved  in  water,  with  the  addition 
of  hydrochloric  acid,  after  which  the  iron  is  oxidized  to  the  ferric 
condition,  unless  it  is  already  wholly  present  in  that  state.  The 
iron  is  then  separated,  by  double  precipitation  with  ammonium 
hydroxide,  as  ferric  hydroxide.  The  precipitate  is  ignited,  and 
weighed  as  ferric  oxide. 

From  the  combined  filtrates  and  washings,  which  must  have 
a  large  volume,  and  which  must  be  free  from  nitrates,  etc.,  the 
sulphate  is  precipitated  by  means  of  a  dilute  solution  of  barium 
chloride.  The  precipitate  is  ignited,  and  weighed  as  barium 
sulphate. 

A.  Procedure  for  the  Determination  of  Iron.  Weigh  out 
into  dry  500  cc.  beakers  two  portions  of  about  i  gram  each, 
and  add  to  each  portion  50  cc.  of  water  and  10  cc.  of  6-normal 
hydrochloric  acid,  keeping  the  beakers  covered  with  watch 
glasses  to  prevent  loss  by  effervescence.  Heat  the  solutions  to 
boihng,  and  treat  each  as  follows:  Add  1 6-normal  nitric  acid 
{before  doing  this,  see  Note  2),  drop  by  drop  with  stirring,  until 
the  darkened  liquid  clears  to  a  yellow,  being  careful  not  to  add 
more  than  i  cc.  altogether.  Boil  for  3  minutes,  add  100  cc.  of 
water,  heat  again  to  boiling,  and  add  rapidly  with  stirring  25  cc. 
of  6-normal  ammonium  hydroxide  solution.  Digest  the  mix- 
ture, which  should  smell  strongly  of  ammonia,  with  heating  and 
occasional  stirring,  until  the  precipitate  is  well  coagulated; 
allow  to  settle ;  and  decant  the  hot  liquid  through  a  9-cm.  ash- 
less filter,  washing  the  precipitate  twice  by  decantation  with 
25  cc.  portions  of  hot  water,  and  leaving  it  mainly  in  the  beaker. 
(Neutralize  the  filtrate  and  washings  with  hydrochloric  acid, 


GRAVIMETRIC  ANALYSIS  6i 

and  begin  the  evaporation  of  the  liquid  at  this  point,  as  directed 
under  the  determination  of  sulphur.    See  p.  4.) 

Dissolve  the  precipitate  by  slowly  pouring  over  the  filter  a 
boiling  mixture  of  5  cc.  of  water  and  10  cc.  of  6-normal  hydro- 
chloric acid,  and  receiving  the  filtrate  (and  washings)  in  the  beaker 
containing  the  bulk  of  the  precipitate,  which  also  will  completely 
dissolve.  After  thoroughly  washing  the  filter,  tear  it  into  small 
bits  and  add  these  to  the  ferric  chloride  solution.  Dilute  to  100 
CO.,  heat  to  boiling,  and  add  as  before  25  cc.  of  6-normal  am- 
monia. Digest  for  some  time,  near  the  boiling  temperature,  and 
decant  the  hot  liquid  through  a  fresh  filter,  washing  the  pre- 
cipitate with  hot  water,  at  first  by  decantation  and  finally  on 
the  filter,  until  5  cc.  of  washings  give  no  cloudiness  with  a  drop 
or  two  each  of  nitric  acid  and  silver  nitrate  solution.  (Neu- 
tralize the  filtrate  and  washings  as  before,  add  the  liquid  to  that 
from  the  first  precipitation,  and  allow  the  evaporation  to  con- 
tinue.) 

Ignite  the  precipitate,  together  with  the  filter  paper,  in  the 
usual  manner  in  an  inclined  platinum  or  porcelain  crucible; 
after  the  decomposition  of  the  paper,  increase  the  flame  to  the 
full,  and,  with  free  access  of  air,  heat  for  20  minutes.  Cool 
in  the  desiccator  and  weigh ;  remember  to  repeat  this  operation 
to  the  attainment  of  constant  weight.  Report  the  percentage 
of  iron  in  the  sample. 

Notes.  —  i.  In  contact  with  air,  the  iron  in  ferrous  salt  solutions  is 
slowly  oxidized  by  dissolved  oxygen  (4  Fe++-h02+2  H20=4  Fe+++ 
-|-4  0H~),  and,  unless  the  solution  contains  free  acid  (which  will  remove 
the  hydroxide  ions),  the  iron  will  be  partially  precipitated  in  the  form  of  a 
basic  ferric  salt.  Moreover,  upon  boiling  an  aqueous  solution  of  ferric 
sulphate,  owing  to  the  increase  of  hydrolysis  with  the  temperature,  there 
results  a  partial  precipitation  of  the  iron,  again  as  a  basic  ferric  salt  {e.g. 
Fe2(S04)3-f-2  H+OH-  ±^  2  Fe(OH)S04+H2S04).  This  action,  however,  is 
prevented  by  the  presence  of  sufficient  hydrochloric  acid ;  the  hydrogen-ion 
of  the  acid,  by  mass  action,  prevents  the  ionization  of  the  water. 

2.  The  complete  oxidation  of  the  iron  is  necessary,  since  ferrous-ion 
is  not  quantitatively  precipitated  by  ammonia;    in  the  absence  of  air, 


62  QUANTITATIVE  CHEMICAL  ANALYSIS 

indeed,  ammonium  salts  are  capable  of  preventing  entirely  the  precipitation 
of  iron  by  ammonia  from  ferrous  salt  solutions.  Ferric-ion,  on  the  other 
hand,  is  completely  precipitated  by  ammonia,  even  in  the  presence  of 
ammonium  salts. 

The  nitric  acid  oxidizes  the  iron  according  to  the  equation : 

6  FeS04+2  HNO3+6  HC1=2  Fe2(S04)3+2  FeCl3+2  NO+4  H2O, 

and  the  darkening  of  the  solution  is  due  to  the  union  of  nitric  oxide  with 
still  unoxidized  ferrous  salt,  to  form  unstable  FeS04(N0) ;  this,  however, 
is  soon  decomposed,  and  the  nitric  oxide  expelled.  (C/.  the  brown-ring 
test  for  nitric  acid.) 

To  insure  the  presence  of  iron  wholly  in  the  ferric  condition,  a  very  small 
quantity  of  the  oxidized  solution  should  be  tested  on  white  porcelain  with 
a  drop  of  freshly  made  potassium  ferricyanide  solution  (dissolve  a  cubic 
miUimeter  of  the  salt  in  15-20  cc.  of  water).  If  the  iron  solution  has  a 
volume  of  50  cc,  and  an  ordinary  drop  a  volume  of  0.05  cc,  then  the  loss 
of  the  latter  would  occasion  an  error  of  o.i  %  in  the  iron  (and  in  the  sulphur) 
determination.  In  this  case,  therefore,  we  should  add  a  couple  of  drops 
of  the  iron  solution  to  i  cc.  of  water  in  a  clean  watch  glass,  and  use  i  drop 
of  this  mixture  for  the  test ;  the  remainder  should  be  washed  back  into  the 
beaker  containing  the  bulk  of  the  solution.  The  error  is  thus  reduced  to 
0.01%,  which  is  negligible. 

Much  time  can  often  be  saved  by  testing  in  this  way  the  solution  made 
from  a  little  of  the  (unweighed)  original  sample,  for  ferrous  iron ;  in  Us 
absence  the  addition  of  the  nitric  acid,  and  its  subsequent  removal  by  evaporation 
to  dryness,  should  be  omitted. 

3.  Upon  the  addition  of  ammonia  in  slight  excess  only,  to  the  ferric  salt 
solution,  the  resulting  precipitate  would  be  sure  to  contain  basic  salts  in 
appreciable  quantity,  and  the  separation  of  the  sulphate  radical  would  be 
incomplete.  By  digesting  the  precipitate  with  a  large  excess  of  ammonium 
hydroxide,  however,  these  basic  compounds  are  more  or  less  completely 
decomposed  into  ferric  hydroxide  and  the  corresponding  ammonium  salts ; 
nevertheless,  since  a  small  quantity  of  the  basic  sulphate  might  escape 
this  action,  the  precipitate  is  redissolved  in  hydrochloric  acid  and  the  opera- 
tion repeated.  In  this  way,  the  quantitative  separation  of  the  sulphate 
radical  is  assured. 

4.  Owing  to  the  solvent  action  of  alkaline  liquids  on  glass,  the  filtration 
should  be  promptly  carried  out,  and  the  filtrate  and  washings  at  once 
neutralized.  The  precipitate,  which  is  of  a  colloidal  nature,  will  give  no 
trouble  if  it  is  filtered  ofif  hot,  and  is  promptly  washed  with  very  hot  water. 
If  allowed  to  cool  off  in  the  mother  liquid,  or  if  washed  with  water  that  i» 


GRAVIMETRIC  ANALYSIS  63 

too  cold,  it  is  almost  sure  to  become  slimy,  and  very  difficult  to  prepare  for 
ignition. 

5.  During  the  decomposition  of  the  paper  and  the  combustion  of  the 
carbon,  the  precipitate  may  be  reduced  in  part  to  Fe304,  and  this  must  be 
oxidized  back  to  Fe203.  Therefore,  during  ignition,  the  air  should  have 
free  access  to  the  oxide;  and,  in  order  to  insure  a  porous  mass  which  is 
readily  reoxidized,  the  filter  originally  used  is  macerated  with  the  solution 
before  the  second  precipitation. 

Concerning  the  effect  of  ammonium  chloride  upon  ferric  and  aluminum 
hydroxides  during  ignition,  see  H.  W.  Daudt,  Jour.  Ind.  Eng.  Chem.,  vol.  7, 
p.  847  (1915)- 

6.  The  foregoing  method  may  be  used  for  the  gravimetric  determination 
of  chromium  or  aluminum,  but  at  the  time  of  filtration  the  solution  should 
contain  but  a  very  sHght  excess  of  ammonia;  chromium  and  aluminum 
hydroxides  are  appreciably  soluble  in  strong  ammonia.  Owing  to  the 
tendency  of  aluminum,  when  thrown  down  by  ammonia,  to  carry  down  and 
retain  basic  compounds,  and  also  to  the  slimy  nature  of  the  precipitate  so 
obtained,  it  is  better  in  the  case  of  this  metal  to  add  sodium  thiosulphate  in 
excess  to  the  dilute  solution,  and  then  to  boil  off  the  sulphur  dioxide,  add 
ammonia  in  slight  excess,  and  filter  boiling  hot.  This  gives  a  compact, 
easily  prepared  precipitate,  which  on  ignition  yields  pure  AI2O3 : 

2  AICI3+6  H0H±^2  Al(OH)3+6  HCl; 
and  6  HCl+3  Na2S203=6  NaCl+3  H2O+3  SO2+3  S. 

If  it  is  desired  by  this  method  to  determine  the  chromium  in  an  alkali 
chromate,  the  latter  is  boiled  with  hydrochloric  acid  and  alcohol,  in  order 
to  reduce  the  chromium  to  the  trivalent  condition: 

K2Cr207+3  C2H6O+8  HC1=2  KCl+2  CrCl3+3  C2H4O+7  H2O. 

While  chromic  hydroxide  is  decomposed  by  heat  into  water  and  the  oxide, 
the  latter  upon  ignition  is  oxidized  to  a  small  extent  to  Cr2(Cr04)3;  in  the 
case  of  this  metal,  therefore,  the  oxide  should  be  ignited  in  a  Rose  crucible, 
in  an  atmosphere  of  hydrogen  (G.  Rothaug,  Zeitschrijt  fiir  anorganische 
Chemie,  vol.  84,  pp.  165-189  (1913)). 

7.  The  hydroxides  of  all  three  metals  are  precipitated  by  sodium  or  potas- 
sium hydroxide,  but  the  precipitates  are  always  contaminated  with  alkaU. 
Furthermore,  aluminum  and  chromium  hydroxides  dissolve  readily  in 
caustic  alkali,  yielding  complex  anions  to  which  the  formulas  A102~  and 
Cr02~  are  usually  ascribed,  but  which  more  likely  are  [A1(0H)4]~  and 
[Cr(0H)4]-.  (C/.  AuCl3+HCl=H[AuCl4];  and  Cr(OH)3+NaOH 
=  Na[Cr(0H)4].) 


64  QUANTITATIVE  CHEMICAL  ANALYSIS 

The  trivalent  ions  of  iron,  chromium,  and  aluminum  form  stable  com- 
plexes with  tartrate  and  citrate  ions  (and  with  certain  other  organic  ma- 
terials) ;  so  that  tartrates,  for  example,  are  capable  of  entirely  preventing 
the  precipitation  of  Fe(0H)3,  Cr(0H)3,  and  A1(0H)3  by  alkali  hydroxides. 
These  and  other  analogous  facts  are  often  made  use  of  in  analytical 
chemistry. 

B,  Procedure  for  the  Determination  of  Sulphur.  Evaporate 
to  dryness,  on  the  steam  bath,  the  combined  filtrates  and  wash- 
ings from  the  iron  determination;  add  lo  cc.  of  6-normal  hy- 
drochloric acid;  and  again  evaporate  to  dryness.^  To  the 
residue  add  loo  cc.  of  water,  heat  to  boiling,  and  filter  into  a 
700  cc.  beaker,  thoroughly  washing  with  hot  water  until  the 
filtrate  and  washings  have  a  volume  of  400  cc.  Heat  this  solu- 
tion to  boiling,  and,  with  stirring,  quickly  pour  in  a  boihng-hot 
mixture  made  by  adding  15  cc.  of  i-normal  barium  chloride 
solution  to  100  cc.  of  water.  Digest  with  stirring  for  a  few 
minutes  at  the  boiling  temperature ;  remove  the  flame  and  allow 
to  stand  for  half  an  hour ;  and,  after  testing  for  complete  pre- 
cipitation, decant  the  clear  Hquid  through  a  filter.  Wash  the 
precipitate  by  decantation,  and  then  upon  the  filter,  with  hot 
water,  until  5  cc.  of  washings  give  no  test  for  chloride-ion.  Dry 
the  filter  and  precipitate  and  ignite  to  constant  weight,  as  in  the 
preceding  determination.  Report  the  percentage  of  SO4  in  the 
sample. 

Notes.  —  i.  Barium  sulphate,  to  a  greater  degree  than  most  precipi- 
tates, has  the  property  of  carrying  down  various  soluble  salts  which  may 
be  present,  and  these  cannot  be  removed  by  simply  washing  the  precipitate. 
This  is  especially  true  of  the  salts  of  trivalent  metals  (as  iron,  chromium, 
etc.),  and  of  nitrates  and  chlorates.  If,  therefore,  nitric  acid  has  been 
used  to  oxidize  the  iron,  it  must  be  completely  removed  by  evaporation  to 
dryness  with  an  excess  of  hydrochloric  acid. 

Iron  is  always  present  in  the  sulphate  precipitated  from  solutions  con- 
taining ferric  salts ;  the  precipitate  evolves  sulphuric  acid  on  heating,  and 

1  In  case  the  original  sample  contained  no  ferrous  iron,  and  the  addition  of  nitric 
acid  was  omitted,  it  is  sufficient  only  to  evaporate  the  neutralized  filtrates  and 
washings  to  about  300  cc.  This  liquid  is  then  transferred  to  the  beaker,  diluted 
to  about  400  cc,  and  treated  further  as  described  in  the  procedure. 


GRAVIMETRIC  ANALYSIS  6$ 

3aelds  low  results  in  spite  of  its  iron  content.  Pure  barium  sulphate  itself 
is  not  decomposed  at  a  red  heat,  but  suffers  loss,  probably  of  sulphur  tri- 
oxide,  at  a  temperature  above  900°. 

The  solubility  of  barium  sulphate  is  about  2.5  parts  per  million  of  water, 
but  it  is  somewhat  more  soluble  in  hydrochloric  acid,  even  very  dilute; 
in  many  salt  solutions  it  is  still  more  soluble. 

2.  In  the  precipitation  of  sulphuric  acid  with  barium  chloride,  the  solu- 
tion should  contain  only  salts  of  the  alkali  metals  and  ammonium,  and  it 
should  be  free  from  nitrates  and  chlorates.  Even  alkali  salts  and  barium 
chloride  are  carried  down  to  some  extent  by  barium  sulphate,  more  or  less 
in  proportion  to  their  concentration,  and  consequently  the  solution  should 
be  dilute.  Since,  further,  the  solubility  of  barium  sulphate,  as  well  as  the 
amount  of  barium  chloride  carried  down,  increases  with  the  concentration 
of  the  hydrochloric  acid  present,  the  quantity  of  the  latter  should  be  re- 
duced to  a  minimum;  some,  however,  must  be  present,  since  otherwise 
the  precipitate  would  be  very  fine  grained,  and  therefore  difficult  to  filter. 

Barium  sulphate  carries  down  quantities  of  chlorine  varying  from  traces 
to  as  much  as  1%,  depending  upon  the  conditions,  and  these  should  there- 
fore be  very  carefully  regulated.  For  a  quantity  of  sulphate  corresponding 
to  1-2  g.  of  BaS04,  the  latter  should  be  precipitated  from  a  solution  which 
has  been  diluted  to  about  400  cc,  and  which  should  contain  1.5  cc.  of  free 
6-normal  hydrochloric  acid.  This  solution  should  be  boiling  hot,  and,  for 
each  gram  of  barium  sulphate,  10  cc.  of  i-normal  barium  chloride  solution 
diluted  to  100  cc,  and  boiling  hot,  should  be  poured  in,  all  at  once,  with  con- 
stant stirring.  In  this  way  exact  results  can  be  obtained,  but  only  owing  to  a 
compensation  of  errors :  although  a  very  small  quantity  of  barium  sulphate 
remains  dissolved  in  the  acidified  salt  solution,  an  approximately  equal 
weight  of  barium  chloride  is  contained  in  the  ignited  precipitate. 

If  precipitated  as  described,  and  with  the  use  of  a  good  filter,  the  pre- 
cipitate shows  no  tendency  to  pass  through  the  pores  of  the  paper;  it  is 
even  possible  to  use  with  success  a  Munroe  crucible  in  this  determination. 

3.  The  barium  sulphate  may  be  partially  reduced  to  the  sulphide  upon 
ignition  with  a  paper  filter ;  in  order  to  prevent  this  as  far  as  possible,  the 
crucible  should  not  be  heated  above  dull  redness  until  the  carbon  has  been 
consumed.  Ignition  in  the  presence  of  air  will  then  suffice  to  oxidize  the 
sulphide  back  to  sulphate. 


66  QUANTITATIVE  CHEMICAL  ANALYSIS 

THE   DETERMINATION  OF  SULPHUR  IN  A   SULPHIDE   ORE 

Method.  The  ore  is  heated  with  strong  nitric  acid  and  potas- 
sium chlorate,  in  order  to  oxidize  the  sulphur  to  sulphuric  acid. 
After  removing  the  nitric  acid  and  chlorate,  as  well  as  the  iron, 
lead,  etc.,  the  sulphuric  acid  is  precipitated  with  barium  chloride, 
and  weighed  as  barium  sulphate. 

Procedure.  Treat  0.25-0.50  g.  samples  of  the  finely  pul- 
verized ore  (depending  upon  the  sulphur  content)  in  250  cc. 
Erlenmeyer  flasks  with  10  cc.  of  nitric  acid  (sp.  gr.  1.42),  and 
heat  very  gently  until  the  red  fumes  have  somewhat  abated. 
Then  increase  the  heat,  and  add  to  the  quietly  boiling  liquid 
potassium  chlorate,  from  time  to  time,  in  o.i  g.  portions,  until 
any  free  sulphur  which  has  separated  is  entirely  oxidized  and 
dissolved ;  finally  add  0.5  g.  of  solid  sodium  chloride  and  evapo- 
rate the  solution  to  dryness.  After  cooling,  cautiously  add 
10  cc.  of  hydrochloric  acid  (sp.  gr.  1.19),  heat  gently  until  solu- 
tion is  as  complete  as  possible,  and  evaporate  to  dryness.  Take 
up  in  5  cc.  of  strong  hydrochloric  acid,  heat  to  boiling,  and 
dilute  with  100  cc.  of  cold  water.  To  the  cold  solution  add  three 
drops  of  methyl  orange,  and  ammonia  to  alkaline  reaction; 
then  add  5  cc.  more  of  ammonia  and  10  cc.  of  ammonium  car- 
bonate solution.  Heat  to  boiling,  allow  the  precipitate  to  settle 
in  the  hot  Hquid,  and  filter  while  still  hot,  washing  thoroughly 
with  hot  water,  and  receiving  the  filtrate  and  washings  in  a  700-cc. 
beaker.  Neutralize  the  filtrate  with  hydrochloric  acid,  and  add 
1.5  cc.  of  the  6-normal  acid  in  excess.  Dilute  the  solution  to 
400  cc,  heat  to  boihng,  and  add  with  stirring  a  boiling-hot  mix- 
ture of  10  cc.  of  I -normal  barium  chloride  solution  and  100  cc. 
of  water.  Allow  the  mixture  to  stand  for  half  an  hour,  test  for 
complete  precipitation,  and  finish  the  determination  as  described 
in  the  previous  Procedure.  Report  the  percentage  of  sulphur 
in  the  sample. 

Notes.  —  i.  Barium  sulphate,  if  present  in  the  ore,  remains  practically- 
unaffected  by  the  above  acid  treatment.    If  it  is  desired  to  determine  the 


GRAVIMETRIC  ANALYSIS  67 

total  sulphur  in  ores  containing  barium,  the  hydrochloric  acid  solution, 
after  the  removal  of  nitrates  and  chlorates  and  dilution  with  100  cc.  of 
water,  may  be  treated  with  5  g.  of  solid  ammonium  chloride  (to  hold  any 
lead  in  solution),  heated  to  boiling,  and  filtered  from  the  insoluble  residue. 
The  filter  containing  the  latter  is  destroyed  by  ignition  in  a  platinum  crucible, 
and  the  residue  fused  with  an  excess  of  sodium  carbonate.  The  fusion  is 
extracted  with  hot  water  and  the  residue  washed  with  sodium  carbonate 
solution;  the  filtrate  and  washings,  which  contain  sodium  sulphate,  are 
added  to  the  hydrochloric  acid  filtrate  containing  the  bulk  of  the  sulphur. 
The  united  filtrates  are  then  treated  with  ammonia  and  ammonium  car- 
bonate, as  described  above. 

2.  The  small  amount  of  sodium  chloride  is  added  before  the  first  evapora- 
tion in  order  to  prevent  the  possible  loss  of  any  free  sulphuric  acid  which 
might  be  present,  —  as,  for  example,  in  the  analysis  of  pyrites.  The  potas- 
sium chlorate  added  would,  however,  probably  be  sufl&cient  in  most  cases 
to  accomplish  this  result.  In  the  analysis  of  pyrites,  which  contains  a 
very  high  percentage  of  sulphur,  samples  should  be  used  of  only  0.25  g. 
Otherwise  the  procedure  is  the  same. 

3.  Upon  adding  ammonia  in  excess  to  the  solution  and  heating  to  boiling, 
there  is  practically  no  danger  of  losing  sulphur  in  the  form  of  basic  ferric 
sulphate.  (In  this  connection  see  Note  3  of  the  procedure  for  iron.)  The 
ammonium  carbonate  is  added  in  order  to  remove  any  lead  which  may  be 
present,  as  the  carbonate,  and  thus  prevent  the  loss  of  sulphur,  as  PbS04, 
before  the  precipitation  with  barium  chloride. 

4.  In  neutralizing  a  solution  with  the  use  of  methyl  orange,  the  solution 
should  be  cold,  since  otherwise  the  methyl  orange  is  not  a  satisfactory 
indicator. 

5.  The  student  should  be  sure  to  read  the  notes  on  the  determination 
of  sulphur  in  iron  sulphate. 


1 


68  QUANTITATIVE  CHEMICAL  ANALYSIS 

THE  DETERMINATION  OF   POTASH  IN    SOLUBLE   SALTS 

The  sample  may  be  a  pure  salt,  a  soluble  industrial  product, 
or  an  artificial  mixture  of  potassium  chloride  and  sodium  car- 
bonate. 

Principle.  The  determination  of  potassium  by  this  method 
depends  upon  the  insolubihty  of  potassium  perchlorate,  and  the 
solubility  of  sodium  and  certain  other  perchlorates  in  96% 
alcohol.  This  is  not  a  precipitation  method,  but  one  of  ex- 
traction.    If  heavy  metals  are  present,  they  are  first  removed. 

Procedure.  Weigh  out  samples  sufficient  to  contain  about 
0.25  g.  of  K2O,  into  50  cc.  beakers,  and  treat  each  as  follows: 
Warm  the  sample  with  25  cc.  of  water,  and,  if  sulphates  are 
absent,  filter  into  a  100  cc.  porcelain  dish ;  if,  however,  sulphates 
are  present,  acidify  with  6-normal  hydrochloric  acid,  stir,  treat 
the  hot  acid  hquid  with  barium  chloride  solution  in  slight  excess, 
filter  into  a  100  cc.  porcelain  dish,  evaporate  the  filtrate  to  dry- 
ness on  the  steam  bath,  and  warm  the  residue  with  15-20  cc.  of 
water,  with  stirring.  Add  to  the  solution  sufficient  perchloric 
acid  to  contain  1.7  times  the  sample's  weight  of  HCIO4  (see  Note 
2),  and  evaporate  to  a  sirupy  consistency.  Add  15  cc.  of  hot 
water  and  2  cc.  of  perchloric  acid,  and  again  evaporate.  Once 
more  add  15  cc.  of  hot  water,  and  evaporate  until  heavy  fumes 
of  perchloric  acid  appear. 

Allow  the  mixture  to  cool  thoroughly,  add  20  cc.  alcohol  con- 
taining 0.2%  by  weight  of  HC104,^  and  stir  for  some  time, 
keeping  the  salt  as  coarsely  granular  as  possible.  Let  settle, 
decant  the  liquid  through  a  weighed  Gooch  crucible^  (containing 
a  mat  moistened  with  the  wash  liquid),  and  to  the  residue  add 
a  second  20-cc.  portion  of  the  wash  liquid.  Stir,  let  settle,  again 
decant,  and  then  drive  ofiF  the  remaining  alcohol  on  the  steam 

*  Made  by  mixing  1.7  cc.  of  the  60%  acid,  or  4.4  cc.  of  the  30%  acid,  with  one 
liter  of  96%  alcohol. 

*  It  is  better  to  use  a  Munroe  crucible,  with  a  filter  of  platinum  sponge.  The 
crucible  itself  may  be  of  gold,  to  save  expense. 


GRAVIMETRIC  ANALYSIS  69 

bath.  Dissolve  the  residue  in  15  cc.  of  hot  water,  add  a  few 
drops  of  perchloric  acid,  and  evaporate  to  heavy  fumes.  Cool, 
add  I  cc.  of  the  wash  liquid,  decant,  and  test  a  few  drops  of  the 
washings  for  complete  extraction.  (The  extraction  with  the 
alcohoHc  liquid  must  be  continued  until  a  few  drops  of  the  filtrate 
leave  no  residue  when  evaporated  to  dryness  on  platinum  foil.) 
Finally,  cool,  add  i  cc.  of  the  wash  liquid,  and  sweep  the  salt 
into  the  Gooch  crucible  with  a  policeman,  washing  at  last  with 
a  very  little  pure  96%  alcohol.  Dry  the  salt  for  half  an  hour 
at  130°,  and  weigh. 

Report  the  percentage  of  K2O  in  the  sample. 

Notes.  —  i.  This  method  was  proposed  in  1831  by  Serullas,  but,  owing 
to  some  mistaken  ideas  concerning  the  properties  of  perchloric  acid,  the 
proposition  did  not  receive  the  attention  it  deserved.  Perchloric  acid 
solutions  of  satisfactory  grade  can  now  be  obtained  in  the  market,  they 
can  be  kept  indefinitely  in  glass-stoppered  bottles,  and  the  method  rivals 
in  results  the  chloroplatinic  acid  process ;  and  this  at  a  greatly  reduced  cost. 

2.  The  specific  gravities  of  perchloric  acid  solutions  are  as  follows: 
7o%HC104,  1.67;  60%  HCIO4,  1.54;  50%  HCIO4,  1.41;  30%  HCIO4, 1.20; 
20%  HCIO4,  1. 1 2.  The  strength  of  a  solution  of  the  pure  acid  may  easily  be 
determined  by  the  dilution  of  a  known  amount  and  titration  with  sodium 
hydroxide,  with  phenolphthalein  as  indicator. 

3.  In  order  to  obtain  the  potassium  as  pure  KCIO4  by  this  method,  it  is 
essential  that  no  strong  acids  be  present,  other  than  perchloric  acid,  which 
yield  salts  insoluble  in  alcohol.  Sodium  chloride  and  sulphate  are  such 
salts,  and  it  is  therefore  necessary  to  remove  chlorides  and  sulphates  before 
the  treatment  with  alcohol.  HCl  may  be  removed  by  repeatedly  evaporat- 
ing the  aqueous  solution  with  the  less  volatile  HCIO4;  but  H2SO4  is  less 
volatile  than  HCIO4,  and  cannot  be  expelled  in  this  way.  Before  the  first 
evaporation,  therefore,  the  latter  should  be  precipitated  from  the  hot  acid 
solution  by  means  of  BaCl2  in  slight  excess.  Phosphates,  though  often 
insoluble  in  alcohol,  need  not  be  removed ;  but,  in  their  presence,  a  larger 
excess  of  HCIO4  should  be  used,  to  insure  their  complete  removal  by  the 
wash  liquid,  as  H3PO4.  (See  Note  5  of  this  procedure,  and  also  Note  6 
under  the  determination  of  calcium.) 

4.  Since  NH4CIO4  is  only  sparingly  soluble  in  alcohol,  ammonium  salts 
should  be  carefully  expelled  by  gentle  ignition,  before  the  treatment  with 
HCIO4.  Moderate  amounts  of  barium,  calcium,  and  magnesium  do  not 
interfere  with  the  procedure ;  their  perchlorates  are  soluble  in  alcohol. 


70  QUANTITATIVE  CHEMICAL  ANALYSIS 

5.  The  perchlorate  mixture  is  extracted  with  alcohol  containing  a  small 
amount  of  HCIO4  because,  owing  to  the  common  ion  efifect,  the  solubility  of 
KCIO4  is  less  in  it  than  in  pure  alcohol ;  the  two  solubiHties  are  about  4  mg. 
and  16  mg.  per  100  cc,  respectively.  The  solubility  of  KCIO4  is  still  less 
in  the  presence  of  sodium,  and  other  soluble  perchlorates ;  ix.  in  the  first 
portions  of  the  alcoholic  extract.  Alkali  and  alkali-earth  phosphates  are 
decomposed  by  perchloric  acid  and  the  H3PO4  dissolves  in  the  acid-alcoholic 
liquid ;  in  the  presence  of  phosphates,  therefore,  perchloric  acid  should  be 
present  in  considerable  excess. 

6.  If  a  known  weight  of  NaCl-KCl  mixture,  obtained  for  example  in  a 
silicate  analysis,  is  converted  into  a  mixture  of  the  perchlorates,  and  the 
KCIO4  isolated  and  weighed,  the  method  yields  both  the  K2O  and  Na20 
contents  of  the  original  sample.     (Cf.  Part  V,  Problem  29.) 

7.  In  order  to  prevent  the  loss  of  KCIO4,  in  the  separation  of  sodium 
and  potassium,  it  has  been  suggested  to  extract  the  perchlorate  mixture 
with  an  alcohoHc  liquid  which  has  previously  been  saturated  with 
KCIO4.  This  procedure,  however,  is  apt  to  lead  to  high  results,  owing  to 
the  precipitation  of  small  amounts  of  potassium  from  the  wash  liquid  by  the 
NaC104  entering  into  solution ;  it  is  therefore  not  to  be  recommended.  In 
order  to  obtain  exact  results,  it  suffices  to  avoid  the  use  of  unnecessary 
quantities  of  the  wash  liquid.  (See,  however,  G.  P.  Baxter  and  M. 
Kobayashi,  Jour.  Amer.  Chem.  Soc,  Vol.  42,  pp.  735  and  2046  (1920).) 


GRAVIMETRIC  ANALYSIS  71 


THE  DETERMINATION  OF  CARBON  DIOXIDE  IN  LIMESTONE 

Method.  The  weighed  carbonate  is  placed  in  an  apparatus 
which  contains  acid  in  a  separate  compartment;  the  whole 
apparatus  is  then  weighed.  After  this  the  acid  is  run  in  upon 
the  carbonate,  and  the  carbon  dioxide  set  free  is  removed  from 
the  apparatus  through  a  tube  filled  with  calcium  chloride,  which 
prevents  the  escape  of  moisture  from  the  apparatus.  Finally, 
the  apparatus  is  weighed  again,  and  the  loss  in  weight  indicates 
the  quantity  of  carbon  dioxide  in  the  sample. 

Many  different  forms  of  apparatus  have  been  devised  for  this 
purpose.  The  one  shown  in  the  accompanying  figure  is  an  im- 
proved form  of  the  so-called  alkalimeter  of  Mohr.  It  consists 
of  a  small,  wide-mouthed,  flat-bottomed  flask  F,  which  has  a 
ground-glass  connection  with  the  tubes  A  and  By  which  are  for 
acid  and  calcium  chloride.  The  ground-glass  joints  are  lubri- 
cated with  a  mixture  of  vaseline  and  beeswax,  or  other  suitable 
substance. 

Procedure.  Thoroughly  clean  the  apparatus,  allow  it  to 
drain,  and  finally  dry  it  by  gently  heating  the  flask  while  draw- 
ing a  current  of  dry  air  through  it,  by  means  of  an  aspirator. 
During  aspiration  the  tubes  C  and  D  should  be  connected  with 
c  and  d,  as  shown  in  the  figure,  to  prevent  the  entrance  of  mois- 
ture into  the  apparatus.  After  drying  the  apparatus,  cover  the 
bottom  of  B  with  a  loose  wad  of  cotton,  and,  through  a  cylinder 
of  glazed  paper  inserted  in  the  neck  of  By  introduce  small  pieces 
of  calcium  chloride  until  the  tube  is  three  fourths  full;  upon 
removing  the  paper  funnel,  be  careful  to  keep  the  upper  walls 
of  the  tube  free  from  the  salt.  Place  another  wad  of  cotton 
above  the  calcium  chloride,  insert  the  stopper,  and  close  the  tube 
at  d  by  slipping  over  it  a  short  rubber  tube  stoppered  with  a 
piece  of  glass  rod.^ 

*  The  tube  must  be  kept  closed  when  not  m  use,  to  prevent  the  gradual  absorp- 
tion of  moisture  from  the  air.  Each  of  the  ordinary  calcium  chloride  tubes  pre- 
viously mentioned  is  filled  in  the  same  way  about  two  thirds  full,  but  in  this  case  a 


72 


QUANTITATIVE  CBEMICAL  ANALYSIS 


GRAVIMETRIC  ANALYSIS  73 

When  all  is  ready,  weigh  out  into  the  flask  about  1.5  g.  of  the 
finely  powdered  substance,  and  add  3-4  cc.  of  water.  Close 
the  stopcock  T  and  fill  the  tube  A  about  three  fourths  full  of 
hydrochloric  acid  (i  volume  of  6-normal  acid  to  1.5  volumes 
of  water)  by  means  of  a  small  funnel.  The  whole  apparatus, 
with  the  tubes  open  at  c  and  d,  is  now  accurately  weighed ;  the 
two  calcium  chloride  tubes  are  connected  at  c  and  d\  and  the 
stopcock  T  is  slightly  opened  so  that  the  acid  from  A  slowly 
drops  into  the  flask.  After  a  quiet  evolution  of  carbon  dioxide 
has  begun,  leave  the  apparatus  to  itself  for  about  half  an  hour, 
until  all  of  the  acid  has  entered  the  flask. 

In  order  wholly  to  remove  the  carbon  dioxide  from  the  ap- 
paratus, connect  the  calcium  chloride  tube  D  with  the  aspirator 
W,  as  shown  in  the  figure  (this  bottle  is  of  course  much  larger 
than  the  figure  would  indicate),  and,  by  means  of  a  screw  clamp, 
regulate  the  flow  of  water  through  e  so  that  air  will  be  drawn 
through  the  flask  F  at  the  rate  of  only  2  or  3  bubbles  per  second. 
Then  heat  F  gently,  by  means  of  a  small  flame,  until  the  acid 
just  begins  to  boil ;  at  once  remove  the  flame,  and  continue  to 
aspirate  air  through  the  apparatus  until  it  is  cold. 

Stopper  the  tubes  at  c  and  d,  wipe  the  apparatus  with  a  clean 
dry  towel,  and  allow  it  to  stand  for  one  half  hour  near  the  bal- 
ance. Finally,  remove  the  stoppers  from  c  and  d,  and  weigh 
the  apparatus.     Report  the  percentage  of  CO2  found. 

Notes.  —  i.  This  method  yields  excellent  results  in  the  estimation  of 
large  amounts  of  carbonic  acid,  in  limestones  and  baking  powders,  but  it 
is  unreliable  for  the  determination  of  small  quantities. 

2.  Since  baking  powders  are  decomposed  by  water,  they  should  be  kept 
dry  until  after  the  apparatus  has  been  weighed ;  and  since  their  efficiency 
as  leaveners  depends  upon  the  volume  of  gas  liberated  under  the  conditions 
of  actual  usage,  water  should  be  placed  in  the  tube  A,  instead  of  acid. 


softened  cork  stopper,  pierced  by  a  short  piece  of  glass  tubing  with  rounded  ends, 
is  introduced  and  shoved  far  into  the  tube  with  the  help  of  a  stirring  rod,  leaving 
the  outer  2  or  3  mm.  empty.  This  space  in  the  tube  is  filled  with  molten  sealing- 
wax,  so  that  an  air-tight  connection  is  made.  These  tubes  are  also  closed,  when 
not  in  use,  by  glass  rods  within  rubber  tubing. 


74  QUANTITATIVE  CHEMICAL  ANALYSIS  ^ 

Otherwise  the  procedure  is  the  same.    The  loss  in  weight  is  then  a  measure 
of  the  available  carbon  dioxide  of  the  sample. 

3.  Carbon  dioxide  is  readily  displaced  from  the  apparatus  by  the  method 
described,  but  in  order  to  insure  its  complete  removal  at  least  a  liter  of  air 
should  be  drawn  through  the  apparatus.  The  small  tube  from  A  should 
project  well  below  the  surface  of  the  liquid,  for  in  order  to  remove  carbon 
dioxide  efficiently  from  the  solution,  the  air  must  be  made  to  bubble  through 
the  liquid. 

4.  Since  commercial  calcium  chloride  is  apt  to  contain  free  lime,  it 
should  for  the  best  results  be  treated  with  carbon  dioxide  before  the  deter- 
mination is  made.  For  this  purpose  a  current  of  the  dry  gas  is  passed 
through  the  apparatus  for  a  minute  or  two,  the  tubes  at  c  and  d  are  closed, 
and  the  apparatus  allowed  to  stand  overnight.  The  carbon  dioxide  is 
then  removed  by  aspirating  dry  air  through  the  apparatus  for  about  20 
minutes,  after  which  the  sample  may  be  placed  in  the  flask. 

5.  The  most  serious  objection  to  this  method  is  the  fact  that,  owing  to 
the  size  and  weight  of  the  apparatus,  there  is  likely  to  be  an  appreciable 
error  in  the  difference  between  the  two  weights.  This  danger,  however, 
can  be  largely  overcome  if  a  similar  piece  of  apparatus  is  available  as  a  tare. 
(See  Part  I,  p.  21.) 


GRAVIMETRIC  ANALYSIS  75 

THE  DETERMINATION  OF   CALCIUM  AND   MAGNESIUM 
OXIDES  IN  LIMESTONE 

Method.  The  hydrochloric  acid  extract  of  the  Hmestone  ^  is 
freed  from  dissolved  silica;  treated  with  bromine  water  and 
ammonia,  to  remove  iron,  aluminum,  manganese,  etc. ;  and  from 
the  filtrate  the  calcium  is  separated  by  double  precipitation 
with  ammonium  oxalate,  and  the  precipitate  ignited  to  the 
oxide.  The  united  filtrates  from  the  calcium  oxalate  are  treated 
with  sodium  phosphate  and  ammonia,  and  the  precipitate  of 
magnesium  ammonium  phosphate  is  ignited  to  magnesium 
pyro-phosphate. 

A.  Procedure  for  the  Determination  of  Calcium.  Weigh  out 
into  two  casseroles  0.5-0.6  g.  samples  of  the  finely  ground  rock, 
and  treat  «ach  as  follows :  Cautiously  moisten  the  powder  with 
5  cc.  of  water,  cover  the  casserole,  add  10  cc.  of  6-normal  hy- 
drochloric acid  in  small  portions,  and  evaporate  to  dryness 
on  the  steam  bath.  Then  add  5  cc.  of  water  and  10  cc.  of  the 
hydrochloric  acid,  evaporate  to  dryness,  and  continue  to  heat 
the  dry  residue  on  the  steam  bath  for  30-45  minutes.  Now  add 
25  cc.  of  water  and  10  cc.  of  the  6-normal  acid,  and  heat  for  10 
minutes;  filter  off  the  residue,  washing  it  with  5  cc.  of  dilute 
hydrochloric  acid  and  then  with  hot  water  until  free  from  chlo- 
rine.    (See  Note  i  in  regard  to  the  insoluble  residue.) 

Add  to  the  filtrate  and  washings  bromine  water,  to  a  dis- 
tinctly yellow  tinge ;  then  boil,  and  add  ammonia  until  its  odor 
just  persists  in  the  solution.  Heat  to  boihng,  and  filter  at  once, 
washing  with  hot  water,  and  allowing  the  washings  to  run  into 
the  filtrate.  Now,  with  the  precipitation  beaker  under  the 
funnel,  pour  over  the  filter  25  cc.  of  hot  hydrochloric  acid  (5  cc. 
of  6-normal  acid,  diluted  to  25  cc.) ;  do  not  bother  if  a  sHght 
residue  remains,  but  wash  the  filter  thoroughly  with  hot  water, 
and  reprecipitate  the  iron,  etc.,  from  the  filtrate  and  washings 
with  bromine  water  and  ammonia,  as  already  described.    Trans- 

*  See  Note  i. 


76  QUANTITATIVE  CHEMICAL  ANALYSIS 

fer  the  precipitate  to  the  filter  previously  used,  and  wash  thor- 
oughly with  hot  water,  adding  the  filtrate  and  washings  to  those 
obtained  in  the  first  operation.  (Concerning  the  precipitate, 
see  Note  i.) 

Evaporate  the  combined  filtrates  and  washings  to  about  200 
cc. ;  if  necessary,  add  ammonia  until  its  odor  is  plainly  per- 
ceptible, and  then  to  the  boiling-hot  liquid  add  ammonium 
oxalate  solution  gradually  with  stirring,  as  long  as  a  precipitate 
continues  to  form,  finally  adding  an  excess  of  10  cc.  Boil  with 
stirring,  for  a  short  time,  set  aside  for  30  minutes,  and  filter  as 
usual;  washing  the  precipitate  twice  by  decantation  with  hot 
water  containing  a  very  little  ammonium  oxalate  and  am- 
monia. (Add  to  the  filtrate  and  washings  5  cc.  of  ammonium 
oxalate  solution,  let  stand  for  10-15  minutes,  and,  if  no  precipi- 
tate forms,  acidify  with  hydrochloric  acid  and  reserve  for  Pro- 
cedure B.) 

Redissolve  the  oxalate  precipitate  by  pouring  warm  3-normal 
hydrochloric  acid  slowly  over  the  filter  and  receiving  it  in  the 
beaker  containing  the  bulk  of  the  material ;  wash  the  filter  with 
water,  and  then  with  very  dilute  ammonia.  Dilute  the  filtrate 
and  washings  to  200-250  cc. ;  heat  to  boiling ;  add  1-2  cc.  of 
ammonium  oxalate  solution,  and  ammonia  to  a  persistent  odor; 
boil  for  a  minute  or  two,  with  stirring,  and  allow  to  stand  as 
before.  Filter  through  the  filter  previously  used,  transfer  the 
precipitate  to  this  filter,  and  wash  thoroughly  with  hot  water 
containing  a  few  drops  of  ammonium  oxalate  solution  and  also 
of  ammonia.  (The  filtrate  and  washings  should  be  acidified  and 
added  to  those  from  the  first  precipitation.  //  not  already 
started,  the  evaporation  of  this  liquid  should  now  be  begun.) 

Gently  ignite  the  dried  precipitate  and  filter,  until  the  latter 
is  consumed ;  then  increase  the  flame  to  the  full,  and  ignite  for 
40-50  minutes.  Finally  heat  the  crucible  for  5  minutes  over 
the  blast  lamp,  cool  in  the  desiccator,  and  weigh;  repeat  this 
treatment  to  constant  weight  (see  Note  7).  Report  the  per- 
centage of  CaO  found. 


GRAVIMETRIC  ANALYSIS  77 

Notes.  —  i.  The  chief  component  of  limestone,  calcium  carbonate,  is 
readily  attacked  by  hydrochloric  acid,  as  also  are  some  of  the  other  com- 
ponents ;  but  few  limestones  are  so  pure  as  to  dissolve  completely  in  hydro- 
chloric acid.  The  residue  may  contain  quartz,  silicates,  pyrites,  or  other 
refractory  materials,  and  carbonaceous  matter  may  also  be  present. 
Furthermore,  the  insoluble  silicates  of  the  residue  are  apt  to  contain  some 
calcium  and  magnesium.  The  thorough  analysis  of  a  limestone  necessi- 
tates the  use  of  an  elaborate  system  of  procedures  and  the  determination  of 
numerous  substances,  but  for  many  technical  purposes  the  analysis  may  be 
confined  to  the  determination  of  the  insoluble  matter  and  silica,  of  the  oxides 
of  iron  and  aluminum  (including  small  quantities  of  the  oxides  of  titanium, 
manganese,  phosphorus,  etc.),  of  calcium  oxide,  and  of  magnesium  oxide. 

Though  its  nature  and  amount  depend  somewhat  upon  the  time  of  action 
and  concentration  of  the  acid  used,  the  insoluble  residue,  if  ignited  and 
weighed,  would  give  a  fair  approximation  of  the  insoluble  matter  and  silica 
of  the  limestone;  and  the  ignited  ammonia  precipitate  would  roughly 
approximate  the  summation  of  the  oxides  of  iron,  aluminum,  etc.,  in  the 
soluble  portion.  Substances,  however,  such  as  hydrous  silicates,  pyrites, 
and  carbonaceous  matter  in  the  insoluble  residue,  and  ferrous  and  manganous 
oxides  (originally  present  as  carbonates)  in  the  soluble  portion,  would  not 
be  correctly  indicated  by  this  method. 

For  a  description  of  the  complete  analysis  of  limestones,  the  student  is 
referred  to  Bulletin  No.  700  of  the  United  States  Geological  Survey,  by  W.  F. 
Hillebrand. 

2.  The  silicates  present  in  the  limestone  are  subject  to  partial  decom- 
position by  the  acid,  and  the  silicic  acid  must  be  converted  into  insoluble 
silica  by  evaporation  and  heating.  The  residue  is  washed  first  with  dilute 
acid  to  prevent  the  separation  on  the  filter  of  basic  salts  of  iron,  aluminum, 
etc.,  owing  to  hydrolytic  action. 

3.  The  bromine  water  is  added  to  oxidize  ferrous  iron,  and  also  man- 
ganese, which  then  precipitates  as  MnO(OH)2.  This  precipitate  should  be 
filtered  ofiF  at  once,  since  the  alkaline  solution  absorbs  carbon  dioxide  from 
the  air,  with  the  consequent  precipitation  of  a  little  calcium  as  the  carbonate. 
Owing  to  this  fact,  and  also  because  the  precipitated  hydroxides  tend  to 
carry  down  the  hydroxides  of  calcium  and  magnesium,  the  precipitate  is 
redissolved  and  again  precipitated,  to  free  it  from  these  metals. 

4.  In  the  separation  of  calcium  from  magnesium  by  means  of  ammonium 
oxalate,  the  calcium  oxalate  tends  to  carry  down  some  magnesium  oxalate, 
probably  as  a  double  salt ;  this  must  be  removed  by  redissolving  the  pre- 
cipitate and  again  throwing  down  the  calcium  from  the  solution  containing 
only  this  small  quantity  of  magnesium.    If  the  quantity  of  magnesium 


78  QUANTITATIVE  CHEMICAL  ANALYSIS 

originally  present  in  the  limestone  is  relatively  very  small,  the  calcium  can 
be  separated  in  one  operation,  from  a  rather  dilute  solution,  by  the  use  of 
ammonium  oxalate  in  excess  of  the  quantity  required  to  convert  both  the 
magnesium  and  the  calcium  into  oxalates.  (In  this  connection,  see 
T.  W.  Richards,  C.  T.  McCaffrey,  and  H.  Bisbee ;  Zeitschriftfur  anorganische 
Chemie,  vol.  28,  p.  71  (1901).) 

5.  A  little  ammonium  oxalate  is  added  in  the  second  precipitation, 
because  an  excess  of  the  reagent  reduces  the  solubility  of  calcium  oxalate, 
and  also  tends  to  hold  magnesium  in  solution  as  a  double  ammonium 
magnesium  oxalate,  which  is  soluble.  For  the  first  reason,  the  pre- 
cipitate is  washed,  not  with  pure  water,  but  with  water  containing  am- 
monium oxalate  and  ammonia.  These  substances  are  volatilized  in  the 
ignition. 

6.  Calcium  oxalate,  though  almost  insoluble  in  water  (5.6  mg.  of  the 
anhydrous  salt  per  liter  of  water),  and  practically  insoluble  in  acetic  acid,  is 
readily  dissolved  by  hydrochloric  acid.  This  behavior  is  due  to  the  fact 
that  oxalic  acid  lies  about  halfway  in  strength  between  acetic  and  hydro- 
chloric acids.  In  acetic  acid  solution,  the  hydrogen-ion  concentration  is 
too  low  to  appreciably  diminish  the  concentration  of  C2O4 —  ions,  and  hence 
no  solvent  action  takes  place.  In  hydrogen  chloride  solution,  however,  the 
high  hydrogen-ion  concentration  causes  the  calcium  oxalate  to  dissolve 
according  to  the  following  scheme : 

CaC204  -^  CsLdOi  :^  Ca++-|-C20r"      ]        f  HC2O4-  or  even 
(solid)  (dissolved)  ^  I^  < 

HCl   :$  CI-  -h      H+    j        I H2C2O4. 

Upon  the  addition  of  a  base  to  this  solution,  hydroxide-ion  removes 
hydrogen-ion  to  form  water,  C2O4 —  is  regenerated,  and  this  unites  with 
Ca+'*'  to  form  CaC204,  which  precipitates. 

7.  Upon  ignition,  calcium  oxalate  becomes  anhydrous  slightly  above 
180°,  and  at  low  redness  it  is  decomposed  into  calcium  carbonate  and 
carbon  monoxide.  Strong  ignition  converts  the  carbonate  into  the  oxide ; 
in  a  platinum  crucible,  this  conversion  may  be  carried  to  completion  over 
a  large  Meker  burner.  In  any  case,  however,  after  the  filter  is  consumed, 
a  good  muffle  furnace,  or  an  electric  crucible  furnace,  is  preferable  for  the 
ignition.    This  is  especially  true  with  porcelain  crucibles. 

Since  calcium  oxide  (quicklime)  absorbs  moisture  and  carbon  dioxide 
from  the  air,  it  should  be  weighed  as  rapidly  as  possible. 

8.  By  burning  off  the  filter  and  then  evaporating  with  2-3  cc.  of  6-normal 
sulphuric  acid,  the  calcium  oxalate  may  be  converted  into  sulphate,  and 
this  may  be  heated  to  constant  weight.    While  this  procedure  is  preferred 


GRAVIMETRIC  ANALYSIS  79 

by  some  analysts,  it  is  nevertheless  disadvantageous,  since  it  involves  danger 
of  loss  by  spattering.  Moreover,  calcium  sulphate  is  more  readily  decom- 
posed upon  ignition  than  barium  sulphate,  and  there  is  some  danger  of  loss 
on  this  account. 

B.  Procedure  for  the  Determination  of  Magnesium.  Evapo- 
rate the  united  filtrates  and  washings  from  A  on  the  steam 
bath  to  incipient  crystallization,  add  1-2  cc.  of  6-normal  hydro- 
chloric acid,  and  then  with  stirring  add  water  cautiously,  a  little 
at  a  time,  until  the  salts  redissolve.  (If  the  acid  solution  con- 
tains solid  matter  at  this  point,  it  should  be  filtered.)  Care- 
fully add  ammonia  just  to  alkaline  reaction  (methyl  orange) ; 
then  sodium  ammonium  phosphate  solution,  very  slowly  with 
stirring,  as  long  as  precipitation  continues,  and  then  10  cc. 
more.  Finally  add  6-normal  ammonia,  one  volume  for  each 
three  of  the  solution,  stir  vigorously  for  10  minutes,  and 
allow  the  mixture  to  stand  for  at  least  6  hours,  —  preferably 
overnight. 

Filter  off  the  precipitate  and  wash  it  with  2.5%  ammonia 
(i  volume  of  6-normal  ammonia  to  3  of  water) ;  then  redis- 
solve the  precipitate  by  slowly  pouring  over  the  filter  10  cc.  of 
warm  2-normal  hydrochloric  acid  and  receiving  the  acid  in  the 
precipitation  beaker.  Wash  the  filter  four  times  with  10  cc. 
portions  of  hot  water.  Add  to  the  filtrate  and  washings  1-2  cc. 
of  sodium  ammonium  phosphate  solution,  and  then,  with  stir- 
ring, add  from  a  pipette  aqueous  ammonia  drop  by  drop  until 
the  odor  of  ammonia  persists  in  the  liquid.  Stir  for  2  minutes, 
add  to  the  mixture  one  third  its  volume  of  6-normal  ammonia 
and  allow  to  stand  for  6  hours. 

Filter  as  usual,  transfer  the  precipitate  to  the  filter  by  means 
of  2.5%  ammonia,  and  continue  to  wash  with  this  liquid  until 
5  cc.  of  washings  give  no  opalescence  with  nitric  acid  and  silver 
nitrate. 

With  free  access  of  air,  ignite  the  dry  precipitate  and  filter, 
increasing  the  heat  only  gradually,  and  not  above  a  dull  red, 
until  the  paper  is  consumed  and  the  precipitate  is  white.     Fi- 


8o  QUANTITATIVE  CHEMICAL  ANALYSIS 

nally,  ignite  to  constant  weight  over  the  blast  lamp  or  in  a  muffle 
furnace.     Report  the  percentage  of  MgO  found. 

Notes.  —  i.  When  washed  with  pure  water,  magnesium  ammonium 
phosphate  is  hydrolyzed  according  to  the  equation, 

MgNH4P04+H20  -^  MgHP04+NH40H. 

The  precipitate  thereby  loses  its  crystalline  form,  and  is  sure  to  give  a  cloudy 
filtrate.  Ammonium  hydroxide  prevents  this  action  by  mass  action,  and 
for  this  reason  the  precipitation  mixture  is  made  to  contain  2.5%  of  ammonia, 
and  the  washing  is  performed  with  ammonia  of  this  concentration. 

2.  In  the  presence  of  ammonium  salts  in  large  quantity,  as  in  the  first 
precipitation,  the  precipitate  is  only  partially  crystalline ;  it  is  apt  to  contain 
the  mixed  phosphate,  Mg  [(NH4)2P04]2,  which  upon  ignition  leaves  mag- 
nesium metaphosphate,  Mg(P03)2;  and,  if  produced  with  a  too  rapid  addi- 
tion of  ammonia,  it  may  also  contain  some  tri-magnesium  phosphate, 
Mg3(P04)2,  which  upon  ignition  remains  unchanged.  This  precipitate, 
therefore,  should  be  purified  by  solution  in  a  little  hydrochloric  acid,  addition 
of  a  small  quantity  of  sodium  ammonium  phosphate,  and  reprecipitation 
from  this  solution  (which  contains  very  little  ammonium  salt)  by  the  gradual 
addition  of  ammonia. 

3.  Upon  ignition,  the  precipitate  is  converted  into  magnesium  pyro- 
phosphate, with  the  evolution  of  ammonia  and  steam : 

2  NH4MgP04.  6  H20=Mg2P207+2  NH3+13  H2O. 

The  precautions  detailed  in  Part  I  should  be  carefully  observed  in  igniting 
this  precipitate.  The  phosphate  is  subject  to  partial  reduction  by  the 
ammonia  or  by  the  filter,  and  also,  if  too  soon  heated  very  strongly,  the 
precipitate  becomes  glazed  over  and  it  is  then  impossible  to  remove  the 
carbon  by  further  heating.  The  precipitate  is  much  more  readily  ignited 
to  whiteness  in  platinum  than  in  porcelain ;  but  in  case  platinum  is  used, 
especial  care  should  be  taken  to  provide  a  plentiful  supply  of  air,  to  pre- 
vent reduction  accompanied  by  injury  to  the  crucible. 

The  most  satisfactory  procedure  is  to  filter  off  the  magnesium  ammonium 
phosphate  in  a  Munroe  crucible  of  platinum  (with  a  platinum  sponge  filter), 
in  which  the  precipitate  can  be  ignited  without  danger  of  loss.  If  a  good 
muffle  furnace  (preferably  electric,  in  which  the  air  is  pure)  is  available,  a 
Gooch  crucible  of  porcelain  may  be  used. 


GRAVIMETRIC  ANALYSIS  8i 

THE    DETERMINATION    OF    PHOSPHORIC    ANHYDRIDE    IN 
PHOSPHATE  ROCK 

Method.  The  finely  ground  mineral  is  heated  with  nitric 
acid,  the  mixture  evaporated  to  dryness,  and  the  residue  ex- 
tracted with  hot  nitric  acid;  the  solution  is  then  filtered  from 
the  insoluble  siHceous  material.  The  filtrate  is  made  almost 
neutral  with  ammonia,  and  is  treated  with  a  solution  of  am- 
monium molybdate,  in  excess,  to  separate  the  phosphoric  acid 
from  calcium,  iron,  aluminum,  etc.  The  precipitated  ammonium 
phospho-molybdate  is  washed  with  acidified  ammonium  nitrate 
solution,  dissolved  in  ammonium  hydroxide,  and  the  phosphoric 
acid  precipitated  with  magnesia  mixture.  The  magnesium  am- 
monium phosphate  is  finally  ignited  to  magnesium  pyrophos- 
phate, which  is  weighed. 

Procedure.  Weigh  out  into  200  cc.  casseroles  two  0.18-0.20 
g.  portions  of  the  powdered  mineral,  and  treat  each  as  follows. 
Add  to  the  sample  15  cc.  of  6-normal  nitric  acid,  cover  the  vessel, 
and  warm  gently  for  a  short  time ;  then  evaporate  to  dryness 
on  the  steam  bath.  Allow  the  dry  residue  to  remain  for  30-40 
minutes  on  the  steam  bath ;  add  25  cc.  of  6-normal  nitric  acid, 
and  heat  for  a  few  moments  to  dissolve  the  soluble  portion. 
Filter,  and  wash  the  siHceous  residue  and  filter  with  small  por- 
tions of  hot  water,  receiving  the  filtrate  and  washings  in  a  400- 
cc.  beaker;  continue  to  wash  as  long  as  the  water  which  runs 
through  will  give  a  turbidity  with  ammonia,  but  add  to  the 
filtrate  all  washings  in  which  such  a  cloud  is  produced.  Cau- 
tiously, and  with  stirring,  add  ammonia  to  the  filtrate  and 
washings  until  the  liquid  remains  barely  turbid;  then  nitric 
acid,  drop  by  drop,  until  the  solution  clears  up  again.  The  vol- 
ume of  the  solution  at  this  point  should  be  not  over  100  cc. 

Heat  the  solution  to  about  70°,  remove  the  burner,  and  add 
80  cc.  of  a  freshly  filtered  solution  of  ammonium  molybdate. 
Digest  for  a  few  hours  at  60-65°,  3-nd  then  decant  the  clear  liquid 
through  a  filter,  washing  the  yellow  precipitate  (kept  in  the 


82  QUANTITATIVE  CHEMICAL  ANALYSIS 

beaker)  with  acid  ammonium  nitrate  solution/  until  5  cc.  of 
washings  give  no  test  for  calcium  with  ammonia  and  ammonium 
oxalate.  (Before  disposing  of  the  filtrate,  digest  it  at  65°  with 
5-10  cc.  more  of  the  molybdate  reagent.  If  no  more  of  the 
yellow  precipitate  is  formed,  the  liquid  should  be  poured  into 
the  bottle  for  "  Molybdate  Residues.'') 

Place  the  beaker  containing  the  yellow  precipitate  under 
the  funnel,  slowly  pour  over  the  filter  four  separate  10  cc. 
portions  of  a  warm  2.5%  solution  of  ammonia,^  and  carefully 
wash  the  filter  five  times  with  10  cc.  portions  of  hot  water. 

To  the  clear  solution  add  hydrochloric  acid,  drop  by  drop, 
with  stirring,  until  the  yellow  cloudiness  produced  disappears 
only  slowly  upon  stirring.  To  this  solution  add  20  cc.  of  mag- 
nesia mixture  from  a  pipette,  at  the  rate  of  about  i  drop  per 
second,  with  vigorous  stirring  (see  Note  7 ;  the  last  10  cc.  may 
be  added  somewhat  faster).  Let  stand  for  15  minutes,  add  15 
cc.  of  15-normal  ammonia,  and  then,  after  a  period  of  2  or  3  hours, 
decant  the  clear  liquid  through  the  original  filter  and  transfer 
the  precipitate  to  the  filter  by  means  of  2.5%  ammonia  water. 
Continue  the  washing  with  this  liquid  until  5  cc.  of  washings, 
after  acidification  with  nitric  acid,  give  no  opalescence  with 
silver  nitrate  solution. 

Dry  the  precipitate  and  filter,  and  ignite  with  free  access  of 
air,  increasing  the  heat  only  gradually,  and  not  above  a  dull 
red,  until  the  paper  is  consumed  and  the  precipitate  white.  Fi- 
nally, ignite  to  constant  weight  over  the  blast  lamp  or  in  a  muffle 
furnace.    Report  the  percentage  of  P2O5  found. 

Notes.  —  i.  The  dehydration  and  removal  of  any  dissolved  silicic  acid 
is  necessary ;  like  tungstic  acid,  molybdic  acid  tends  to  form  with  silicic  acid 
substances  of  great  complexity,  and  these,  if  present  in  the  phospho-molyb- 
date  precipitate,  would  be  likely  to  interfere  with  its  complete  solubility  in 
ammonia. 

1  Made  by  mixing  50  cc.  of  6-normal  ammonium  hydroxide  with  100  cc.  of 
6-normal  nitric  acid,  and  diluting  the  mixture  with  350  cc.  of  water. 

2  Made  by  mixing  25  cc.  of  6-normal  ammonium  hydroxide  with  75  cc.  of  hot 
water. 


GRAVIMETRIC  ANALYSIS  83 

Calcium  phosphate,  though  almost  msoluble  in  water,  dissolves  readily 
in  nitric  acid,  according  to  the  following  system  of  equilibria : 

Cae(P04)2$Ca3(P04)2$3Ca+++2P04— I  ^HPO";  HzPOr; 
{solid)  {dissolved)  \  or  even  H3PO4. 

:cHN03:^yN03-+y  H+     J 

Therefore,  in  order  to  test  for  phosphate  in  the  acid  washings  from  the  sili- 
ceous residue,  it  is  only  necessary  to  add  ammonium  hydroxide ;  hydroxide- 
ion  removes  hydrogen-ion,  PO4  is  regenerated,  and  it  unites  with  Ca+"*' 
which  also  is  present  to  form  Ca3(P04)2,  which  precipitates. 

Also,  later  on,  when  it  is  desired  to  practically  neutralize  this  solution 
with  ammonium  hydroxide,  it  is  only  necessary  to  add  the  latter  to  the 
appearance  of  a  permanent  turbidity;  this  indicates  the  almost  complete 
removal  of  hydrogen-ion  from  the  solution. 

2.  Nitric  acid,  rather  than  hydrochloric,  is  used  in  the  precipitation  of 
the  phospho-molybdate,  because  molybdates  are  subject  to  a  partial  reduc- 
tion by  hydrochloric  acid.  Although  nitric  acid  alone  would  tend  to  dissolve 
this  precipitate,  this  action  is  counteracted  by  the  presence  of  the  ammonium 
nitrate  which  is  formed  upon  almost  neutralizing  the  original  solution  of 
the  rock  with  ammonium  hydroxide.  Owing  to  this  fact,  in  washing  the 
precipitate  free  from  metals  with  dilute  nitric  acid,  the  wash  liquid  also  is 
made  to  contain  an  equivalent  amount  of  ammonium  nitrate.  It  should  be 
noted,  too,  that  the  molybdate  reagent  contains  ammonium  nitrate  and  free 
nitric  acid.     (See  the  Preparation  of  Reagents,  in  the  Appendix.) 

3.  The  precipitation  of  the  phosphoric  acid  as  magnesium  ammonium 
phosphate  from  the  original  solution  of  the  rock  is  not  possible,  owing  to  the 
presence  of  metals  such  as  iron,  aluminum,  calcium,  etc.,  which  also  form 
phosphates  insoluble  in  ammonia.  Therefore,  the  phosphoric  acid  is  pre- 
viously separated  from  the  metals,  in  acid  solution,  as  ammonium  phospho- 
molybdate. 

4.  While  the  precipitation  of  the  phospho-molybdate  should  be  carried 
out  in  a  warm  solution,  the  maximum  temperature  which  can  be  safely 
employed  is  65° ;  warmer  solutions  tend  also  to  deposit  free  M0O3,  which  is 
white. 

A  large  excess  of  the  molybdate  reagent  is  required  to  effect  a  complete 
precipitation  of  the  phosphoric  acid.  Theoretically  1.95  g.  of  M0O3  are 
required  to  combine  with  the  phosphate  in  0.2  g.  of  rock  containing  40  per 
cent  of  P2O6 ;  while  the  quantity  of  the  reagent  actually  used  (80  cc.)  con- 
tains about  5  g.  of  M0O3.  The  presence  of  ammonium  nitrate  in  the  solution 
is  also  conducive  to  complete  precipitation.    These  substances,  by  mass 


84  QUANTITATIVE  CHEMICAL  ANALYSIS 

action,  prevent  the  partial  dissociation  of  the  complex  into  its  more  soluble 
constituents. 

5.  While  the  composition  of  the  yellow  precipitate  varies  somewhat  with 
the  conditions,  it  nevertheless  seems  to  correspond  pretty  closely  to  the 
formula  (NH4)3P04 .  12  M0O3 .  2  HNO3 .  H2O;  at  any  rate,  under  proper 
analytical  conditions,  the  ratio  P2O6 :  24  M0O3  holds  very  closely.  The  yel- 
low precipitate  dissolves  readily  in  aqueous  ammonia  to  give  ammonium 
phosphate  and  ammonium  molybdate ;  and  molybdate-ion  is  not  precip- 
itated by  magnesia  mixture. 

6.  The  precipitate  obtained  upon  the  addition  of  the  magnesia  mixture 
should  have  the  formula  Mg(NH4)P04 .  6H2O;  in  which  case  it  will  be 
wholly  crystalline.  The  gradual  addition  of  the  reagent,  with  constant  stir- 
ring, is  conducive  to  this  end.  If  magnesia  mixture  is  added  too  rapidly, 
or  without  vigorous  stirring,  it  will  accumulate  locally  in  the  liquid  and  tend 
to  give  a  precipitate  contaminated  with  molybdic  acid  and  magnesia.  Such 
a  precipitate  will  not  be  wholly  crystalline;  it  should  be  redissolved  in  a 
small  volume  of  hydrochloric  acid,  the  solution  treated  with  1-2  cc.  of 
magnesia  mixture,  and  the  hot  liquid  slowly  neutralized  with  2.5%  ammonia. 
Strong  ammonia  is  then  added,  and  the  analysis  continued  as  above. 

Magnesia  mixture  is  prepared  by  putting  together  in  solution  magnesium 
chloride,  ammonium  chloride,  and  ammonium  hydroxide.  The  function 
of  the  ammonium  chloride  is  to  drive  back  the  ionization  of  the  ammonium 
hydroxide,  and  thus  to  prevent  the  precipitation  of  magnesium  hydroxide, 
which  otherwise  might  contaminate  the  precipitate. 

7.  If  the  magnesium  ammonium  phosphate  were  washed  with  pure 
water,  it  would  be  hydrolyzed  according  to  the  equation, 

MgNH4P04+H20  %  MgHP04+NH40H; 

it  would  lose  its  crystalline  form,  and  be  almost  sure  to  give  a  cloudy  filtrate. 
Ammonium  hydroxide,  however,  prevents  this  decomposition  by  mass 
action. 

8.  Upon  the  ignition  of  this  precipitate,  steam  and  ammonia  are  evolved, 
and  magnesium  pyrophosphate  is  formed : 

2  NH4MgP04 .  6  H20=Mg2P207+2  NH3+13  H2O. 

Concerning  the  precautions  to  be  observed  in  the  ignition,  see  Note  3 
under  the  determination  of  magnesium  in  limestone. 


GRAVIMETRIC  ANALYSIS  85 

THE  DETERMINATION  OF  SILICA  IN  A  REFRACTORY 

SILICATE 

Method.  The  powdered  sample,  though  not  readily  attacked 
by  hydrochloric  acid,  is  decomposed  by  molten  sodium  car- 
bonate, with  which  it  yields  sodium  silicate  and  other  com- 
pounds readily  attacked  by  hydrochloric  acid.  By  evaporation 
and  extraction  with  this  acid,  and  ignition  of  the  insoluble  resi- 
due, the  silica  (somewhat  contaminated  with  iron  oxide  and 
alumina)  is  obtained.  After  weighing,  this  is  evaporated  with 
hydrofluoric  and  sulphuric  acids,  and  the  sulphur  trioxide  ex- 
pelled from  the  iron  and  aluminum  sulphates  by  strong  ignition. 
The  weight  of  the  impure  siHca  less  that  of  the  ignited  impuri- 
ties gives  the  weight  of  siHca  originally  in  the  sample. 

Procedure.  Weigh  out  into  platinum  crucibles  two  0.75  g. 
portions  of  the  very  finely  ground  material.  Also  weigh  out 
roughly  two  4  g.  portions  of  pure  anhydrous  sodium  carbonate. 
In  each  case,  place  about  3  g.  of  the  flux  upon  the  sample  in  the 
crucible,  and,  with  the  latter  upon  a  sheet  of  glazed  paper,  mix 
its  contents  with  a  clean  dry  rod.  Stir  the  remaining  gram  of 
flux  with  the  rod,  to  remove  from  the  latter  any  particles  of  the 
mixture,  and  then  pour  it  upon  the  mixture  in  the  crucible. 

With  the  crucible  covered,  slowly  heat  the  mixture  up  to  the 
highest  possibility  of  an  ordinary  burner,  and  then,  if  necessary 
in  order  to  secure  complete  fusion,  heat  the  mixture  over  a  Meker 
burner  or  a  blast  lamp.  As  soon  as  the  mass  is  in  quiet  fusion, 
with  practically  no  further  evolution  of  gas,  take  up  the  crucible 
in  tongs  apphed  to  the  upper  edge,  and,  by  means  of  a  slow 
rotary  motion,  cause  the  molten  substance  to  spread  around  the 
walls  of  the  crucible,  where  it  will  solidify.  Then,  with  the  mass 
still  at  a  dull-red  heat,  lower  the  crucible  for  a  few  seconds  half- 
way down  into  cold  water,  but  with  care  that  no  water  shall 
enter  the  crucible.  Allow  a  Httle  time  for  the  melt  to  reheat 
the  crucible,  and  lower  it  again  into  the  water;  then  set  it 
aside  to  cool.    The  melt  may  then  nearly  always  be  removed 


86  QUANTITATIVE  CHEMICAL  ANALYSIS 

from  the  inverted  crucible  by  gentle  tapping.  {Do  not  deform 
the  crucible) 

Place  the  melt  in  a  300  cc.  beaker,  add  100  cc.  of  water  and, 
with  stirring  and  care  to  avoid  loss,  gradually  add  30  cc.  of  6- 
normal  hydrochloric  acid.  (Also  clean  the  crucible  and  cover 
with  acid,  and  wash  this  liquid  into  the  beaker.  In  case  the 
melt  cannot  be  removed  from  the  crucible,  place  both  in  the 
beaker  and  treat  with  water  and  acid  as  described.)  Heat  the 
beaker  gently,  and,  if  necessary,  aid  the  disintegration  of  the 
melt  by  gentle  pressure  with  the  broadened  end  of  a  glass  rod. 

After  complete  disintegration,  transfer  the  mixture  to  a  cas- 
serole, and  evaporate  on  the  steam  bath,  stirring  towards  the 
end,  until  the  residue  is  dry ;  continue  to  heat  this  on  the  steam 
bath  for  about  an  hour.  Thoroughly  wet  the  residue  with  5  cc. 
of  i2-normal  hydrochloric  acid,  warm  gently,  add  100  cc.  of 
water,  and  heat  to  boiling.  Filter  at  once,  and  wash  with  sev- 
eral small  portions  of  hot  2-normal  hydrochloric  acid,  and  then 
with  hot  water  until  free  from  chlorides;  receive  the  filtrate 
and  washings  in  a  casserole.  Evaporate  to  dryness  on  the 
steam  bath,  heat  the  dry  residue  as  before,  and  repeat  the  ex- 
traction and  filtration,  washing  with  hot  acid  and  then  with  hot 
water,  —  hut  use  a  fresh  filter  to  receive  the  second  yield  of  silica. 

Cautiously  ignite  both  filters,  with  their  contents,  in  a  weighed 
platinum  crucible  (if  only  silica  is  to  be  determined,  it  is  not 
necessary  to  know  the  weight  of  the  crucible),  and,  after  the 
filters  are  consumed,  heat  for  half  an  hour  at  the  full  heat  of  the 
burner,  and  allow  to  cool.  Moisten  the  residue  with  2  or  3 
drops  of  strong  sulphuric  acid,  cautiously  expel  this,  ignite  to 
low  redness,  and  then  for  half  an  hour  over  the  blast  lamp. 
Repeat  the  blasting  for  periods  of  5  minutes,  to  constant  weight. 

Now  add  to  the  silica  in  the  crucible  i  cc.  of  6-normal  sulphuric 
acid  and  3-5  cc.  of  pure  48%  hydrofluoric  acid  {Caution:  this 
acid  is  capable  of  causing  serious  injury  to  the  skin).  Evapo- 
rate as  far  as  possible  on  the  steam  bath  in  a  well-drawing  hood, 
adding  more  hydrofluoric  acid,  if  necessary.    Cautiously  expel 


GRAVIMETRIC  ANALYSIS  87 

the  sulphuric  acid,  heat  to  low  redness,  and  finally  ignite  over 
the  blast  lamp,  for  5-minute  periods,  to  constant  weight.  The 
loss  in  weight  due  to  this  treatment  represents  the  weight  of 
pure  siHca  in  the  sample.    Report  the  percentage  of  Si02  found. 

Notes.  —  i.  Upon  fusion  with  sodium  carbonate,  silicates  are  decom- 
posed with  the  evolution  of  carbon  dioxide,  and  formation  of  products  such 
as  sodium  siHcate  and  aluminate,  ferrous  carbonate,  ferric  oxide,  calcium 
and  magnesium  carbonates,  etc. ;  e.  g.  in  the  case  of  a  certain  species  of 
garnet,  we  may  have 

Mg2FeAl2Si30i2+7  Na2C03=2  MgC03+FeC03H-2  NaA102 
+3  Na4Si04+4  CO2. 

Owing  to  the  evolution  of  gas,  the  heating  'should  be  gradual  and  the 
crucible  should  be  kept  covered. 

2.  Unless  the  sample  is  ground  very  fine,  the  coarser  particles  will  bfe 
sure  to  escape  the  action  of  the  flux ;  the  sample,  therefore,  should  finally 
be  ground  in  an  agate  mortar  until  it  will  wholly  pass  through  a  loo-mesh 
sieve. 

In  case  a  gritty  residue  should  remain  after  the  disintegration  with 
hydrochloric  acid,  which  can  be  detected  by  means  of  a  glass  rod,  the  sample 
has  not  been  completely  decomposed ;  the  analyst  should  start  over  with  a 
fresh  sample. 

3.  Upon  treating  the  fusion  with  considerable  dilute  acid,  sihcic  acid 
at  first  tends  to  dissolve ;  upon  evaporation,  however,  it  is  largely  dehydrated 
and  rendered  insoluble.  The  use  of  concentrated  acid  instead  of  dilute 
would  cause  the  separation  of  a  jelly-like  mass,  in  which  metallic  salts  would 
be  protected  from  the  solvent  action  of  the  acid. 

4.  The  silicic  acid  cannot  be  rendered  wholly  insoluble  by  a  single 
evaporation  and  heating;  nor  are  repeated  evaporations  and  moistenings 
before  filtration  as  effective  in  separating  the  silica  as  are  alternate  evapora- 
tions and  filtrations.  The  second  portion  of  silica  obtained  should  be 
filtered  off  upon  a  fresh  filter ;  otherwise  some  of  the  first  portion  is  apt  to 
reenter  the  solution.    These  facts  have  not  yet  been  satisfactorily  explained. 

5.  To  free  the  silica  as  far  as  possible  from  mineral  salts,  the  residue  after 
evaporation  should  be  thoroughly  extracted  with  warm  hydrochloric  acid ; 
and  the  solution  should  be  diluted  to  a  large  volume  to  prevent  the  inclosure 
of  impurities  by  the  silica.  The  silica  is  first  washed  with  dilute  acid,  to 
prevent  the  partial  separation  of  basic  salts  of  iron,  aluminum,  etc.,  by 
hydrolysis ;  the  washing  is  then  completed  with  hot  water. 


88  QUANTITATIVE  CHEMICAL  ANALYSIS 

6.  The  finely  divided  silica  very  stubbornly  retains  moisture,  and  a 
prolonged  ignition  is  necessary.  Even  then  the  ignited  powder  should  be 
weighed  rapidly,  since  it  tends  to  reabsorb  moisture. 

7.  In  spite  of  every  care,  the  ignited  silica  is  rarely  pure.  Upon  evapora- 
tion with  hydrofluoric  and  sulphuric  acids,  however,  the  silica  is  volatilized 
as  silicon  tetrafluoride  and  steam,  and  a  sulphate  residue  is  left.  If  the 
contaminating  substance  is  an  alkali  salt,  as  sodium  chloride,  the  residue 
will  remain  as  sulphate,  even  at  high  temperatures;  but  certain  other 
sulphates,  as  those  of  iron,  aluminum,  and  titanium,  evolve  sulphur  trioxide 
on  ignition  and  leave  the  corresponding  oxides.  In  the  estimation  of 
silica,  the  weight  of  impurities  in  the  silica  is  always  determined  by  weighing 
the  residue  from  the  hydrofluoric  and  sulphuric  acid  treatment;  in  order 
then  that  the  impurities  weighed  with  the  silica  may  be  as  nearly  as  possible 
identical  with  the  residue  from  this  treatment,  it  is  best  to  treat  the  silica 
before  ignition  with  a  few  drops  of  sulphuric  acid.  The  final  residue  from 
the  hydrofluoric  and  sulphuric  acids  should  also  be  subjected  to  the  same 
temperature  employed  in  the  ignition  of  the  silica. 

8.  The  procedure  for  the  determination,  in  the  united  filtrates  from 
the  siHca,  of  the  mixed  oxides  of  iron,  aluminum,  etc.  (in  which  the  hydroxide 
precipitate  would  be  ignited  together  with  the  residue  from  the  hydrofluoric 
and  sulphuric  acid  treatment),  of  calcium  oxide,  and  of  magnesium  oxide, 
does  not  differ  materially  from  that  given  under  the  determination  of 
calcium  and  magnesium  in  limestones. 

9.  For  a  thorough  study  of  the  analysis  of  silicate  and  carbonate  rocks, 
the  student  is  referred  to  Bulletin  No.  700  oj  the  United  States  Geological 
Survey,  by  W.  F.  Hillebrand. 


GRAVIMETRIC  ANALYSIS  89 

THE  ELECTROLYTIC  DETERMINATION  OF  COPPER 

The  sample  to  be  analyzed  may  be  pure  copper  sulphate,  an 
artificial  mixture  of  the  carbonates  of  copper  and  sodium,  a 
copper  ore,  or  a  nickel  coin.  In  case  stationary  electrodes  are 
employed,  the  solution  should  contain  not  over  0.2  g.  of  copper 
and  5  cc.  of  nitric  acid  (sp.  gr.  1.42),  and  should  have  a  volume 
of  100  cc. ;  in  the  case  of  a  rotating  anode,  however,  the  solution 
may  contain  as  much  as  0.5  g.  of  copper,  and  it  should  contain, 
in  a  volume  of  100  cc,  3-5  cc.  of  nitric  acid  (sp.  gr.  1.42),  or  i  cc. 
of  sulphuric  acid  (sp.  gr.  1.84). 

Method.  The  copper  salt  is  decomposed  by  the  electric 
current,  and  the  copper  deposited  upon  the  cathode  (negative 
electrode).  The  cathode  is  weighed  before  and  after  the  opera- 
tion, and  the  increase  in  weight  indicates  the  quantity  of  copper 
in  the  sample. 

The  polarity  of  the  terminals  may  be  determined  by  bringing 
the  wires,  about  0.5  cm.  apart,  into  contact  with  a  piece  of  filter 
paper  moistened  with  potassium  iodide  solution.  At  the  positive 
terminal  iodine  will  separate  and  color  the  paper. 

Cleaning  the  Platinum  Electrodes}  The  electrodes  are  freed 
from  grease  by  heating  with  dilute  sodium  hydroxide  solution, 
after  which  they  are  washed  with  water;  the  cathode  is  then 
dried,  allowed  to  cool  in  a  desiccator,  and  weighed.  To  clean 
the  platinum  cathode  after  the  determination,  cover  the  deposit 
completely  with  6-normal  nitric  acid,  heat  for  at  least  15  minutes, 
and  wash. 

A.  Procedure  with  Stationary  Electrodes.^  Dissolve  a  0.5- 
0.6-g.  sample  of  copper  sulphate,  CUSO4 .  5  H2O,  in  50  cc.  of 
water,  in  a  tall  150  cc.  beaker;  stir  to  complete  solution,  add 
4  cc.  of  nitric  acid  (sp.  gr.,  1.42),  and  dilute  to  100  cc.  Im- 
merse the  electrodes  in  the  solution  and  connect  them  in  such 
a  way  that  the  electrode  with  the  larger  surface  is  made  the 

*  In  case  a  silver  cathode  is  used,  see  Note  12. 

*  If  the  sample  is  an  ore,  see  Note  11. 


90  QUANTITATIVE  CHEMICAL  ANALYSIS 

cathode.  The  electrolysis,  which  should  be  carried  out  at  a 
potential  of  1.9-2.0  volts,  may  be  completed  in  the  cold  over- 
night, or  in  two  or  three  hours  if  the  temperature  is  kept  at 
70-80°  by  means  of  a  heated  sheet  of  wire  gauze  placed  a  short 
distance  below  the  beaker.  Finally  test  for  complete  deposition 
by  adding  a  Uttle  water  to  raise  the  level  of  the  solution  on  the 
cathode;  if  after  30  minutes  no  copper  is  to  be  seen  upon  the 
fresh  platinum  surface,  the  deposition  is  probably  complete. 
(Test  a  few  cubic  centimeters  of  the  solution  with  sodium  acetate 
and  a  drop  of  potassium  ferrocyanide  solution.)  Without  dis- 
connecting the  electrodes,  siphon  off  the  electrolyte  while  in- 
troducing distilled  water,  until  the  current  ceases  to  pass ;  this 
is  to  prevent  the  re-solution  of  any  of  the  copper  by  the  acid 
liquid.  Remove  the  cathode,  wash  it  with  water,  then  with 
alcohol,  and  dry  it  for  a  short  time  in  an  air  bath  at  85-90°. 
Allow  the  cathode  to  cool  in  a  desiccator,  and  weigh. 

B,  Procedure  with  a  Rotating  Anode. ^  Heat  a  five-cent  coin 
with  sodium  hydroxide  solution  to  free  it  from  grease,  then  wash 
it  with  water,  and  dry  at  100°.  After  cooHng,  weigh  the  coin, 
and  dissolve  it  in  50  cc.  of  6-normal  nitric  acid,  in  a  covered 
casserole.  Evaporate  the  solution  to  dryness  on  the  steam 
bath,  dissolve  the  residue  in  about  100  cc.  of  cold  water,  add 
10  cc.  of  sulphuric  acid  (sp.  gr.,  1.84),  allow  to  cool,  and  transfer 
the  whole  to  a  500  cc.  measuring  flask,  diluting  to  the  mark  with 
water.  Measure  off  one  tenth  of  the  well-mixed  solution  into 
a  50-cc.  graduated  flask,  for  transference  later  to  the  electrolytic 
vessel. 

With  the  anode  attached  to  the  shaft  of  the  rotator  (which  is 
connected  by  means  of  a  mercury  cup,  or  otherwise,  with  the 
positive  terminal),  and  the  cathode  connected  with  the  negative 
terminal,  pour  100  cc.  of  water  into  the  electrolytic  vessel  and 
adjust  the  electrodes  so  that,  when  four  fifths  covered  with 
water,  they  do  not  come  into  contact  with  one  another,  nor 
cause  a  loss  of  liquid,  when  the  anode  is  rotated.  Throw  in 
*  If  the  sample  is  an  ore,  see  Note  11. 


GRAVIMETRIC  ANALYSIS  91 

the  maximum  resistance  of  the  rheostat,  see  that  all  connec- 
tions are  well  made,  and  close  the  switch;  then  draw  off  the 
water,  transfer  the  solution  quantitatively  to  the  vessel  (with 
care  that  the  volume  of  solution  and  washings  is  close  to  100  cc), 
start  the  motor,  and  immediately  decrease  the  resistance  of  the 
rheostat  until  the  ammeter  registers  about  i  ampere.  In  the 
presence  of  nickel^  the  voltage  should  not  exceed  2.7.  After  about 
55  minutes,  test  for  complete  deposition  by  adding  a  Uttle  water 
to  raise  the  level  of  the  solution  on  the  cathode;  if  after  10 
minutes  no  copper  is  visible  on  the  freshly  exposed  platinum,  the 
deposition  is  complete.  When  this  is  the  case,  without  discon- 
necting the  terminals,  stop  the  rotator  and  draw  off  the  solution 
into  a  large  beaker,  carefully  pouring  in  water  as  fast  as  the  solu- 
tion flows  out.^  As  soon  as  the  ammeter  indicates  that  no  cur- 
rent is  passing,  throw  off  the  switch,  remove  the  cathode  and 
wash  off  the  water  with  a  little  alcohol ;  dry  below  icx)°,  allow 
to  cool  in  a  desiccator,  and  weigh. 

Notes.  —  i.  If  two  platinum  plates,  immersed  in  an  aqueous  solution 
of  copper  sulphate,  are  connected  by  wire  with  the  poles  of  a  storage  battery, 
metallic  copper  will  be  deposited  upon  one  of  the  plates;  under  certain 
conditions,  all  of  the  copper  will  separate  in  the  form  of  a  compact,  firmly 
adherent  metallic  film. 

The  process  of  decomposition  is  called  electrolysis;  the  solution  under- 
going decomposition  is  called  an  electrolyte;  the  two  poles  by  which  the 
current  enters  and  leaves  the  electrolyte  are  called  electrodes.  When  salt 
solutions  are  electrolyzed,  the  positive  ions  (cations)  move  towards  the 
negative  electrode  (cathode),  and  the  negative  ions  (anions)  towards  the 
positive  electrode  (anode). 


1  If  it  is  desired  to  determine  the  nickel  electrolytically,  evaporate  this  dilute 
solution  to  a  volume  of  25-30  cc,  make  slightly  alkaline  with  ammonia,  filtering 
off  any  ferric  hydroxide  which  may  be  precipitated,  and  to  the  solution  (40  cc.  in 
volume)  in  the  electrolytic  vessel  add  60  cc.  of  15-normal  ammonia.  Electrolyze 
at  3.0-3.5  volts  with  a  rotating  anode.  After  about  an  hour,  test  for  complete 
deposition  by  adding  to  a  few  drops  of  the  solution,  neutralized  with  acetic  acid, 
a  drop  or  two  of  dimethyl-glyoxime  solution  (a  red  color  indicates  nickel).  When 
the  deposition  is  complete,  proceed  as  directed  in  the  copper  determination. 
Finally  remove  the  nickel  from  the  platinum  cathode  by  heating  for  at  least  15 
minutes  with  6-normal  nitric  acid. 


92  QUANTITATIVE  CHEMICAL  ANALYSIS 

The  quantity  of  electricity  which  passes  through  the  solution  in  unit 
time,  or  the  speed  of  the  current,  is  measured  by  an  ammeter.  The  unit, 
called  an  ampere,  is  represented  by  the  unvarying  current  which,  when 
passed  through  a  solution  of  silver  nitrate,  deposits  metallic  silver  at  the 
rate  of  o.oomS  g.  per  second.^ 

The  electromotive  force,  i.e,  the  electrical  pressure  which  drives  the 
current  along  the  circuit,  is  measured  by  the  voltmeter.  The  unit,  called 
a  volt,  is  represented  by  the  electrical  pressure  that  produces  a  current  of 
one  ampere  when  steadily  applied  to  a  conductor  whose  resistance  is  one 
ohm. 

The  unit  of  resistance,  called  an  ohm,  is  represented  by  the  resistance 
offered  to  an  unvarying  electric  current  by  a  column  of  mercury  14.4521  g. 
in  mass,  of  a  constant  sectional  area  and  a  length  of  106.3  cm.,  at  the  tem- 
perature of  melting  ice. 

These  magnitudes  are  always  related  to  one  another  as  follows  (Ohm's 
law): 

Quantity  of  electricity  (amperes)  =  Electromotive  force  (volts)  ^^  .  ^  |. 

Resistance  (ohms)  R 

A  most  satisfactory  source  of  current  for  electro-analysis,  in  which  a 
steady,  non-fluctuating  current  is  desired,  is  the  lead  storage  cell,  the 
E.  M.  F.  of  which  is  slightly  over  2  volts.  The  voltage  and  other  conditions 
demanded  in  the  work  may  be  provided,  of  course,  by  means  of  a  suitable 
combination  of  such  cells  into  a  battery.  In  practice,  the  potential  dif- 
ference between  the  electrodes  is  regulated  by  means  of  incandescent  lamps, 


coils  of  wire,  or  other  devices,  which  offer  resistance  to  the  flow  of  electricity 
along  the  circuit  and  convert  electrical  energy  into  heat.    Any  variety  of 


*  The  quantity  of  a  given  metal  deposited  by  a  current  of  electricity  is  directly 
proportional  to  the  quantity  of  electricity  which  passes  through  the  solution; 
and  the  quantities  of  different  metals  deposited  by  a  specific  quantity  of  electricity 
are  directly  proportional  to  the  chemical  equivalents  of  the  metals.  These  two 
statements  are  known  as  Faraday's  laws,  though  these  apply  to  nonmetallic  ions 
as  well. 


GRAVIMETRIC  ANALYSIS  93 

rheostat,  with  a  suitable  range  in  resistance,  may  be  used  in  this  work, 
but  the  sliding-contact  coil  resistances  in  the  market  are  especially  con- 
venient. Ammeters  and  voltmeters  should  be  such  as  are  designed  for  a 
limited  range,  so  that  they  may  be  read  with  ease  and  accuracy.  The 
manner  of  connecting  the  instruments  is  illustrated  in  the  figure  on  the 
preceding  page. 

2.  The  passage  of  an  electric  current  of  suitable  voltage  through  the 
solution  of  an  ionogen  is  associated  with  physical  and  chemical  changes 
which  often  may  be  utilized  in  exact  gravimetric  analysis. 

The  chemical  effect  at  the  cathode  is  always  some  form  of  reduction. 
Simple  metallic  ions,  as  those  of  copper,  tin,  nickel,  cobalt,  cadmium,  etc., 
travel  towards  the  cathode,  where  they  give  up  their  charges  and  separate 
in  the  metallic  condition;  while  the  hydrogen  ion  here  loses  its  positive 
charge  and  either  acts  directly  as  a  reducing  agent  (e.g.  nitric  acid  to  am- 
monia) or  is  evolved  as  gaseous  hydrogen. 

At  the  anode,  on  the  other  hand,  the  chemical  effect  is  always  some  form 
of  oxidation.  The  anions  of  the  halogen  group  are  liberated  as  free  chlorine, 
bromine,  or  iodine  and  may  act  as  oxidizing  agents,  while  from  solutions 
containing  hydroxide,  sulphate,  or  nitrate  ions,  oxygen  separates  at  the 
anode  and  either  acts  directly  as  an  oxidizing  agent  or  is  evolved  in  gaseous 
form.  (It  should  be  borne  in  mind  in  this  connection  that  aqueous  solutions 
always  contain  the  ions  of  water.)  Although  positive  ions  always  move 
towards  the  cathode,  certain  metals  (e.g.  lead,  cobalt,  nickel,  and  a  few 
others)  may,  under  specific  conditions,  be  oxidized  (possibly  to  complex 
oxy-anions)  and  deposited  more  or  less  completely  at  the  anode  in  the  form 
of  insoluble  peroxides.  In  fact,  lead  can  be  determined  accurately  in  this 
way,  as  an  oxide. 

3.  Metals,  like  all  other  substances,  possess  when  immersed  in  water  a 
characteristic  solution  tension,  by  which  is  understood  an  expansive  force 
which  seeks  to  drive  particles  of  the  metal  into  the  solution ;  when  a  metal 
is  immersed  in  the  solution  of  one  of  its  salts  it  will  either  send  more  of  its 
atoms  into  the  solution  as  ions,  or  some  of  its  ions  will  be  discharged  from 
the  solution  on  its  surface  as  atoms.  In  the  first  case  the  metal  will  be- 
come negatively  charged,  and  in  the  second  case  positively  charged  with 
respect  to  the  solution;  in  either  case  equilibrium  will  be  reached  when 
the  solution  tension  of  the  metal  is  exactly  counterbalanced  by  the  electro- 
static charges  and  the  osmotic  pressure  of  the  metallic  ions  in  the  solu- 
tion. 

Upon  comparing  the  different  elements  from  the  standpoint  of  the 
potential  difference  between  them  and  their  salt  solutions,  at  identical 
normal  ion-concentrations,  a  characteristic  series  of  values  is  obtained. 


94  QUANTITATIVE  CHEMICAL  ANALYSIS 

In  the  case  of  solutions  of  normal  ion-concentration,  for  example,  some  of 
the  values  are  as  follows : 

Zn= +0.493  Ag=— 1.05 

Cd=  +0.143  1=  —0.80 

Fe= +0.067  Br=— 1.27 

Co=  —0.045  0=  —1.50    (At  N.  H+-ion  concentration.) 

Ni  =  —  0.049  CI  =  —  1 .63 

H= -0.277  S04=-2.i8 
Cu=  —0.606 

4.  In  an  electrolysis,  each  electrode  is  soon  covered  with  a  deposit ;  i.e.  it 
becomes  an  electrode  of  the  material  discharged.  Thus,  a  system  like  that 
discussed  above  comes  into  existence  at  each  electrode;  and  further  elec- 
trolysis is  opposed  by  the  solution  tensions  of  the  discharged  products. 
(The  electrodes  are  said  to  be  polarized.)  In  order  then  to  decompose  a 
salt  solution  continuously,  a  voltage  at  least  slightly  in  excess  of  the  polariza- 
tion voltage  must  be  applied ;  i.e.  a  voltage  greater  than  the  numerical  dif- 
ference between  the  individual  potential  differences  at  the  two  electrodes. 
Assuming  normal  ion-concentrations,  the  decomposition  voltage  of  copper 
chloride,  for  example,  is  —1.63  minus  —0.606=1.02  volts;  but  the  decom- 
position voltage  of  a  solution  as  calculated  in  this  manner,  especially  in  the 
case  of  a  salt  of  an  oxyacid,  frequently  fails  to  agree  with  that  found  by 
experiment.  The  separation  of  gases  at  the  electrodes  is  often  accompanied 
by  "overvoltages,"  which  vary  more  or  less  markedly  with  the  material 
and  physical  nature  of  the  electrodes;  moreover,  the  ion-concentration  is 
generally  unknown,  and  it  always  varies  during  the  electrolysis.  At  any 
rate,  in  the  case  of  any  ionogen  there  always  exists  (under  specific  conditions) 
a  definite  decomposition  voltage,  below  which  the  ionogen  cannot  be  de- 
composed by  the  electric  current. 

In  the  case  of  an  ionogen  of  known  decomposition  voltage,  we  should 
simply  use  a  somewhat  higher  voltage,  but  if  other  metallic  ions  were 
present  it  might  be  impossible  to  completely  deposit  one  metal  without 
using  a  voltage  that  would  start  the  deposition  of  the  second  metal  also. 
While  copper  can  readily  be  separated  from  cobalt  or  from  nickel  by  elec- 
trolysis, it  is  not  possible  to  separate  nickel  from  cobalt  in  this  way ;  and  in 
general  only  metals  whose  deposition  voltages  dififer  by  several  tenths  of  a 
unit  can  be  separated  from  each  other  by  maintaining  an  intermediate 
voltage  during  the  electrolysis. 

The  addition  of  certain  reagents,  as  ammonia,  potassium  cyanide,  am- 
monium oxalate,  etc.,  to  solutions  contkining  two  metals  sometimes  reduces 
the  concentration  of  one  metallic  ion  very  much  more  than  that  of  the 


GRAVIMETRIC  ANALYSIS  95 

other,  owing  to  the  formation  of  more  or  less  stable  complexes,  and  makes 
it  possible  to  perform  a  separation  by  the  "constant  voltage"  method  that 
otherwise  might  not  be  possible. 

5.  The  current  strength  will  of  course  depend  upon  the  voltage  used, 

since,  according  to  Ohm's  law,  i  =  — .       In  performing  an  electrolysis,  the 

R 

voltage  actually  available  is  diminished  by  the  decomposition  voltage  of 

the  electrolyte  (polarization  voltage);    hence  the  current  which  passes  is 

equal  to  the  available  voltage  minus  the  decomposition  voltage  of  the 

electrolyte,  divided  by  the  resistance  of  the  circuit. 

6.  The  quantity  of  metal  deposited  in  a  given  time  is  dependent  upon 
the  strength  of  the  current  in  amperes.  A  current  of  i  ampere  is  capable 
of  depositing  1.118  mg.  of  silver,  and,  according  to  Faraday's  law,  equiva- 
lent amounts  of  other  elements,  per  second.  This  law  might  be  used  to 
calculate  the  time  necessary  for  the  complete  deposition  of  the  metal  if, 
under  the  analytical  conditions,  it  were  the  only  cation  taking  part  in  the 
electrolysis.  In  the  neighborhood  of  the  cathode,  however,  the  concentra- 
tion of  the  solution  with  respect  to  this  cation  gradually  decreases  to  an 
infinitesimal  value,  and  the  resistance  and  the  decomposition  voltage  of  the 
solution  therefore  rise ;  finally  a  point  is  reached  at  which  other  ions  begin 
to  be  discharged.  Since  circulation  of  the  solution  tends  to  maintain  a 
uniform  distribution  of  the  ions,  mechanical  stirring  favors  the  rapid  deposi- 
tion of  those  ions  which  have  the  lowest  discharge  voltages. 

7.  Unless  the  solution  is  mechanically  stirred,  the  rate  of  deposition  of 
a  given  metal  decreases  rapidly,  owing  to  the  decreasing  concentration  of 
its  ions  around  the  cathode  and  to  the  continually  increasing  proportion 
of  the  current  which  is  carried  by  the  hydrogen  (or  other)  ions.  Since  a 
rapid  circulation  of  the  solution  tends  greatly  to  prevent  the  local  decrease 
in  metallic  ion-concentration  around  the  cathode,  and  since  with  improved 
circulation  currents  of  much  higher  density  may  be  used  than  would  other- 
wise give  satisfactory  deposits,  it  is  possible  to  greatly  reduce  the  time  neces- 
sary for  a  determination  by  performing  the  electrolysis  with  the  use  of  a 
rotating  electrode. 

8.  Owing  to  the  reduced  viscosity  at  higher  temperatures,  the  resistance 
offered  by  an  aqueous  solution  to  the  passage  of  electricity  decreases  with 
a  rise  in  temperature,  and  in  this  way  the  voltage  required  to  produce  a 
given  current  may  be  reduced  to  a  minimum;  this  may  sometimes  be 
of  importance  in  electrolytic  separations.  Moreover,  in  case  stationary 
electrodes  are  used,  heating  the  solution  during  electrolysis  gives  rise  to 
more  or  less  rapid  convection  currents,  and  also  increases  the  speed  of 
diffusion,  and  these  effects  are  equivalent  to  a  gentle  mechanical  stirring. 


96  QUANTITATIVE  CHEMICAL  ANALYSIS 

The  solution  should  never  be  heated  to  the  boiling  point,  however,  since  the 
deposit  might  in  that  case  be  loosened  from  the  cathode. 

9.  The  deposited  metal  tends  to  redissolve  in  the  electrolyte  (cf.  "polari- 
zation"), and  consequently  the  rate  at  which  the  metal  is  deposited  must 
exceed  that  at  which  it  redissolves.  The  metal  is  only  deposited  from  the 
solution  in  immediate  contact  with  the  cathode,  so  that  the  greater  the  area 
of  the  cathode,  the  more  metal  there  is  available  for  deposition ;  but  also 
the  greater  the  rate  of  re-solution.  Hence  it  follows  that  the  current 
strength  necessary  for  the  satisfactory  deposition  of  the  metal  is  propor- 
tional to  the  area  of  the  cathode. 

The  current  strength  per  unit  area  is  called  the  current  density;  a  square 
decimeter  is  generally  taken  as  the  unit  area.  Hence  a  "normal  current 
density  of  2  amperes  "  means  a  current  of  2  amperes  per  100  sq.  cm.  of  cathode 
area,  or  of  i  ampere  for  50  sq.  cm.  of  cathode  area,  etc.  While  the  tendency 
of  a  metal  to  redissolve  fixes  a  lower  limit  for  the  current  density  to  be  used, 
a  higher  limit  is  set  by  the  tendency  of  the  metal  to  form  spongy,  non- 
adherent films  when  deposited  too  rapidly. 

It  is  highly  important  for  accurate  work  to  deposit  the  metal  in  the 
form  of  a  compact  film  which  can  easily  be  washed  and  weighed  without 
loss.  The  condition  of  the  deposit  depends  not  only  upon  the  current 
density  used,  but  also  upon  the  concentration  of  the  metallic  ions  in  the 
solution,  the  amount  of  free  acid  and  other  substances  present,  the  tem- 
perature, etc.  The  best  conditions  for  specific  cases  have  been  determined 
by  repeated  experiments. 

10.  Concerning  the  effect  upon  the  nature  of  the  deposit  of  the  products 
that  accumulate  in  the  solution  during  an  electrolysis,  that  are  purposely 
added  to  the  solution,  or  which  were  originally  present  in  the  sample,  it 
may  be  stated  that  a  very  marked  influence  is  often  exerted  by  certain 
acids,  bases,  and  other  substances.  A  solution  of  copper  sulphate,  if 
electrolyzed  without  the  addition  of  another  substance,  is  almost  sure  to 
give  a  reddish  brown,  non-adherent  deposit  of  spongy  copper;  the  addi- 
tion of  a  little  sulphuric  acid  gives  rise  to  a  much  more  compact  deposit, 
while  the  addition  of  nitric  acid  leads  to  a  still  better  deposit  of  bright  red 
firmly  adherent  metal.  A  small  quantity  of  urea,  in  addition  to  either  acid, 
appears  to  favor  still  more  the  formation  of  a  satisfactory  deposit.  On  the 
other  hand,  high  current  densities,  which  cause  a  rapid  discharge  of  hydro- 
gen, are  apt  to  yield  loosely  adherent  deposits  of  spongy  metal.  A  current 
density  which  gives  a  bright  red,  coherent  deposit  of  pure  copper  when  no 
interfering  substance  is  present,  will  often  give  a  very  dark,  loosely  adherent 
deposit  when  arsenic  is  present,  even  in  small  amount.  Such  impurities 
must  be  removed  before  the  electrolysis  is  begun. 


GRAVIMETRIC  ANALYSIS  97 

11.  If  the  sample  to  be  analyzed  by  this  method  is  a  copper  ore,  and  is 
not  known  to  be  free  from  arsenic  and  other  interfering  substances,  it  should 
be  subjected  to  special  treatment,  in  order  to  obtain  a  solution  suitable  for 
electrolysis.  In  most  cases,  a  satisfactory  solution  may  be  prepared  accord- 
ing to  the  procedure  detailed  under  the  volumetric  estimation  of  copper 
{which  see) ;  the  ore  is  evaporated  with  aqua  regia,  the  residue  extracted 
with  dilute  hydrochloric  acid  and  water,  and  the  copper,  arsenic,  etc.,  pre- 
cipitated with  sodium  thiosulphate ;  the  arsenic  is  then  driven  off  by  igni- 
tion, the  residue  evaporated  to  dryness  with  nitric  acid,  and  finally  taken 
up  in  4  cc.  of  nitric  acid  (sp.  gr.,  1.42)  and  50  cc.  of  water.  This  solution 
is  diluted  to  100  cc.  and  electrolyzed. 

12.  In  electro-analysis,  owing  to  its  resistance  to  attack  by  the  electro- 
lytic solutions,  platinum  has  been  the  material  most  often  used  for  electrodes. 
The  present-day  price  of  this  metal,  however,  tends  greatly  to  encourage 
the  use  of  other  materials.  Other  metals,  as  silver  and  copper,  are  in  some 
cases  suitable  for  use  as  cathodes  in  the  deposition  of  metals  (as  is  also  the 
more  expensive  palladium-gold  alloy  which  is  in  the  market),  but  the  anode 
must  still  be  made  of  platinum  or  of  something  equally  resistant.  In  the 
determination  of  copper,  for  example,  a  silver  cathode  is  just  as  satis- 
factory as  one  of  platinum ;  the  deposit  should  be  removed  in  the  cold  by 
means  of  dilute  hydrochloric  acid,  with  the  addition  of  a  little  hydrogen 
peroxide  or  nitric  acid,  and,  after  washing  with  ammonia,  the  cathode  is 
again  ready  for  use. 

Since  an  upper  limit  is  set  for  the  current  density,  the  time  necessary 
for  the  deposition  of  a  metal  decreases  with  the  increasing  area  of  the 
cathode;  for  this  reason  the  cathode  should  offer  a  large  surface  for  the 
deposit.  A  thin  platinum  dish  has  a  comparatively  large  inner  surface,  and 
is  especially  satisfactory  for  certain  determinations.  A  platinum  disk  or  a 
flat  spiral  of  platinum  wire  may  be  used  as  the  other  electrode.  The  elec- 
trodes commonly  used,  however,  are  more  economical.  They  are  hollow 
cylinders  of  thin  foil  or  of  fine  mesh  gauze,  and  elongated  spirals  of  heavy 
platinum  wire ;  the  cyHnders,  used  to  receive  the  deposit,  if  of  platinum,  weigh 
10-12  g.,  and  the  wire  spirals  about  8  g.  The  gauze  cylinders  are  the  most 
efl5cient ;  they  present  a  larger  surface,  all  parts  of  which  equally  receive 
the  deposit,  and  they  allow  a  free  circulation  of  the  electrolyte  and  con- 
sequently the  use  of  higher  current  densities. 

These  cylindrical  cathodes  and  spiral  anodes  are  the  ones  assumed  in 
the  foregoing  procedures.  The  use  of  silver  cathodes  is  advocated  in  the 
determination  of  cooper,  but  they  should  not  he  heated  with  the  hydrochloric 
acid-hydrogen  peroxide  cleaning  mixture,  owing  to  the  solvent  action  in  that 
case. 


98  QUANTITATIVE  CHEMICAL  ANALYSIS 

13.  In  case  the  electrolysis  is  to  be  performed  with  stationary  electrodes, 
it  is  best  to  use  a  tall  beaker  of  small  diameter,  which  can  be  heated.  If 
heating  is  not  desired,  however,  or  if  a  rotating  electrode  is  used,  the  most 
suitable  vessel  is  a  150-cc.  glass  cylinder,  with  a  rounded  bottom  ending  in 
an  outlet  tube  provided  near  the  top  with  a  stopcock ;  the  electrodes  should 
reach  nearly  to  the  bottom  of  this  cylinder,  to  insure  efficient  mixing.  After 
the  deposition  is  complete,  without  interrupting  the  current,  the  electrolyte 
can  easily  be  drawn  off  with  the  simultaneous  introduction  of  distilled  water 
above. 


PART   III 

VOLUMETRIC  ANALYSIS 
GENERAL  DISCUSSION 

Fundamental  Principles.  In  volumetric  analysis,  the  quan- 
tity of  an  element  or  compound  present  in  a  weighed  sample 
is  determined  by  calculation  from  the  volume  of  a  dissolved  re- 
agent of  accurately  known  concentration  which  is  required 
to  complete  a  definite  reaction.  The  concentration  or  value  of 
the  solution  for  producing  a  given  reaction  is  determined  by  an 
operation  known  as  standardization;  the  solution  is  caused  to 
react,  under  favorable  conditions,  with  a  known  weight  of  a 
definite  substance  and,  from  the  volume  of  the  solution  re- 
quired, its  reaction  value  may  readily  be  calculated.  This  solu- 
tion is  then  called  a  standard  solution. 

The  value  of  standard  solutions  may  be  expressed  in  terms 
of  the  weight  of  reagent  actually  present  in  each  cubic  centimeter, 
or,  better,  in  terms  of  the  weight  of  a  given  substance  with  which 
one  cubic  centimeter  of  the  solution  will  react.  But  since  the 
weight  of  reagent  present  in  a  unit-volume  is  always  chemically 
equivalent  to  the  weight  of  substance  with  which  the  unit- volume 
reacts,  it  is  in  general  more  convenient  to  express  the  value  of 
the  standard  solution  in  terms  of  chemical  equivalents  per  unit- 
volume  ;  i.e.  in  terms  of  a  normal  solution.  Such  solutions,  for 
example,  may  be  made  up  to  half-normal  tenth-normal,  or 
fiftieth-normal  concentration.^ 

'  The  value  of  the  standard  solution  to  be  used  in  a  specific  case  will  depend 
largely  upon  the  weight  of  the  sample  involved  in  the  titration.  In  the  estimation 
of  acids  and  bases  in  usual  amounts,  it  is  customary  to  use  half-normal  solutions ; 

99 


100  QUANTITATIVE  CHEMICAL  ANALYSIS 

A  normal  solution  contains  in  one  liter  one  gram-equivalent 
of  the  active  reagent;  i.e,  that  quantity  which  is  chemically 
equivalent  to  8.000  g.  of  oxygen,  or  to  1.008  g.  of  hydrogen. 
Thus,  a  normal  acid  solution  contains  i  .008  g.  of  available  hydro- 
gen-ion per  liter  {e.g.  one  mol  of  HCl,  or  one  half  mol  of  H2SO4) ; 
and  a  normal  alkali  solution  contains  in  one  liter  sufficient  avail- 
able hydroxide-ion  to  combine  with  1.008  g.  of  hydrogen-ion, 
or  17.008  g.  {e.g.  one  mol  of  NaOH,  one  half  mol  of  Ba(0H)2,  etc.). 
A  normal  oxidizing  solution  has  an  oxidizing  value  per  liter  of 
8.000  g.  of  oxygen  {e.g.  one  gram-atom  of  available  iodine,  one 
sixth  mol  of  K2Cr207,  etc.) ;  and  a  normal  reducing  solution  has 
a  reducing  value  per  liter  of  i  .008  g.  of  hydrogen  {e.g.  one  mol  of 
FeS04,  one  half  mol  of  SnCl2,  etc.).  It  will  be  seen  that  (if  they 
react  quantitatively)  equal  volumes  of  normal  acid  and  alkali 
will  exactly  neutralize  one  another,  equal  volumes  of  normal  oxi- 
dizing and  reducing  reagents  will  exactly  use  up  one  another,  etc. 

Nevertheless,  it  must  be  realized  that  the  equivalent  or  normal 
weight  of  a  reagent  may  not  always  be  the  same.  Thus  the 
normal  weight  of  oxahc  acid  is  one  half  its  molecular  weight, 
whether  used  as  an  acid,  as  a  reducing  agent,  or  as  a  precipitating 
agent ;  but  the  normal  weight  of  nitrous  acid  would  be  its  molec- 
ular weight  if  it  were  used  as  an  acid,  or  to  oxidize  hydriodic 
acid,  while  if  used  to  reduce  potassium  permanganate  it  would 
be  only  half  as  large.  In  the  case  of  potassium  permanganate, 
one  molecule  yields  5  equivalents  of  available  oxygen  in  the  ti- 
tration of  ferrous  iron,  but  only  3  in  the  titration  of  manganese ; 
the  normal  weight  of  this  oxidizing  agent  may  therefore  be  either 
one  fifth  or  one  third  of  the  molecular  weight,  depending  upon  the 
reaction. 

Although  for  special  purposes  it  is  sometimes  worth  while  to 
prepare  exactly  normal  or  tenth-normal  solutions,  it  is  in  general 
much  easier,  and  just  as  satisfactory,  to  prepare  and  use  stand- 
but  in  the  determination  of  small  amounts,  tenth-normal  solutions  give  more 
accurate  values,  since  their  state  of  dilution  minimizes  the  effect  of  errors  of  measure- 
ment. 


VOLUMETRIC  ANALYSIS  loi 

ard  solutions  which  closely  approximate  these  in  value.  In 
practice,  the  volumes  used  of  such  solutions  can  readily  be  ex- 
pressed in  terms  of  those  of  exactly  normal  (or,  for  that  matter, 
tenth-normal)  concentration ;  e.g.  25.75  cc  of  a  0.0987  N  solution 
are  equivalent  to  25.75X0.0987=2.542  cc.  of  the  normal,  or 
to  2.542X10=25.42  cc.  of  the  tenth-normal  solution. 

The  Ftmdamental  Reactions  of  Volumetric  Analysis.  Volu- 
metric analytical  processes  are  usually  based  upon  sharply  defined 
chemical  reactions  which  can  be  made  to  run  rapidly  to  comple- 
tion in  aqueous  solution.  This,  however,  is  not  always  effected 
by  direct  titration,  i.e.  by  simply  bringing  the  reacting  sub- 
stances together  in  equivalent  quantities.  In  many  cases  a 
standard  solution  is  first  added  in  appreciable  excess,  and  the 
excess  then  titrated  with  a  second  standard  solution.  In  other 
cases  the  substance  is  treated  with  some  compound  with  which 
it  reacts  to  set  free,  or  to  carry  down  as  an  insoluble  precipitate, 
a  definite  proportion  of  some  other  substance ;  and  the  latter  is 
subsequently  titrated  with  a  standard  solution. 

According  to  the  nature  of  the  titration  reactions  themselves, 
volumetric  analysis  is  subdivided  as  follows : 

A.  Neutralization  Methods;  utilized  for  the  determination 
of  acids  and  bases,  and  therefore  called  methods  of  acidimetry 
and  alkalimetry. 

B.  Methods  of  Oxidation  and  Reduction;  of  which  the  most 
important  are  the  permanganate,  and  iodometric  processes  — 
e.g.  the  determination  of  ferrous  iron  or  of  manganese  by  means 
of  potassium  permanganate. 

C.  Precipitation  Methods;  as,  for  example,  the  determination 
of  silver  by  means  of  sodium  chloride  solution. 

The  reactions  employed  in  connection  with  these  processes 
are  in  general  of  the  reversible  type.  Of  these,  the  reactions  of 
neutralization  run  to  completion  in  consequence  of  the  low  de- 
gree of  ionization  of  one  product,  water ;  those  of  oxidation  and 
reduction,  in  consequence  of  the  relative  potentials  of  the  oxidiz- 
ing and  reducing  substances  under  specific  experimental  condi- 


I02  QUANTITATIVE  CHEMICAL  ANALYSIS 

tions ;  and  the  reactions  of  precipitation,  owing  to  the  formation 
and  precipitation  of  substances  of  an  insoluble  nature. 

Determination  of  the  End-Point.  In  order  to  utilize  a  re- 
action for  purposes  of  titration,  it  is  necessary  to  have  a  means 
of  recognizing  the  point  at  which  an  equivalent  quantity  of  the 
standard  solution  has  been  added.  In  approaching  or  passing 
this  point  certain  properties  of  the  solution  are  subject  to  rapid 
change ;  there  may  be  a  sudden  change  in  the  color,  the  electri- 
cal conductivity,  or  the  oxidation  potential  of  the  solution,  or  a 
precipitate  may  cease  to  form.  In  numerous  instances  the 
presence  of  a  substance  purposely  added,  called  an  indicator, 
may  give  rise  to  a  well  defined  color  change  in  the  solution,  or  it 
may  cause  a  precipitate  to  form  and  thus  render  the  solution 
cloudy  at  the  end.  The  point  to  which  the  standard  solution 
must  be  added  in  order  to  make  such  a  change  apparent  is  called 
the  end-point  of  the  titration.  The  process  is  the  more  satis- 
factory the  smaller  the  difference  between  this  point  and  the 
point  at  which  an  exact  equivalent  of  the  titration  reagent  has 
been  added;  in  most  of  the  estabHshed  volumetric  processes, 
this  difference  is  exceedingly  small. 

In  a  few  cases,  in  which  there  is  an  appreciable  difference  be- 
tween the  true  and  the  observed  end-point  of  a  reaction,  it  is 
possible,  by  taking  advantage  of  the  principle  of  compensating 
errors,  to  obtain  very  accurate  results;  any  errors  which  are 
involved  in  the  actual  analysis  are  offset  by  equal  errors  made 
in  the  standardization.  In  general,  the  difference  between  the 
true  and  the  observed  end-point  will  remain  constant  under  the 
same  specific  conditions;  if,  therefore,  the  standardization  and 
the  analysis  are  both  performed  under  identical  conditions,  the 
standard  solution  simply  serves  as  a  means  of  comparing  two 
approximately  equal  quantities  of  the  same  substance,  one  of 
which  is  accurately  known.  From  the  data  of  standardization, 
for  example,  we  calculate  the  value  of  the  solution  per  cubic 
centimeter  in  terms  of  the  substance;  and  then,  by  multiply- 
ing the  number  of  cubic  centimeters  consumed  in  the  titration 


VOLUMETRIC  ANALYSIS  103 

by  this  value,  we  obtain  the  weight  of  the  substance  contained 
in  the  sample  analyzed. 

Bearing  this  in  mind,  it  will  be  seen  that,  at  least  from  a 
theoretical  standpoint,  the  highest  attainable  accuracy  cannot 
be  expected  in  a  given  case  unless  the  standard  solution  used  has 
had  its  value  determined  in  terms  of  the  particular  substance 
under  investigation,  and  under  conditions  in  every  way  identi- 
cal with  those  prevalent  throughout  the  analysis. 

General  Theory  of  Indicators.  When  a  substance  is  titrated 
in  the  presence  of  an  indicator,  the  end-point  is  recognized  as 
the  visible  effect  of  a  chemical  change  in  which  the  indicator  it- 
self takes  an  active  part.  Either  the  indicator  may  combine 
with  the  reagent  in  the  standard  solution  (when  the  latter  has 
been  added  in  slight  excess),  or  it  may  enter  into  combination 
with  a  portion  of  the  substance  undergoing  titration,  to  be  again 
liberated  in  colorless  form  upon  the  addition  of  the  standard 
solution.  These  two  cases  are  exemplified  in  the  behavior  of 
starch  as  an  indicator,  in  the  processes  of  iodometry;  in  those 
titrations  which  are  made  with  iodine,  a  blue  color  appears 
and  persists  upon  the  addition  of  the  slightest  excess  of  iodine, 
while  in  titrations  of  iodine  with  sodium  thiosulphate  solution, 
for  example,  the  blue  color  of  the  iodo-starch  is  discharged  upon 
the  addition  of  the  thiosulphate. 

In  those  cases  in  which  the  reagent  acts  upon  the  indicator, 
the  reactions  involved  in  the  titration  may  be  represented  by 
the  following  equations,  in  which  S  is  the  substance  titrated, 
R  the  reagent  in  the  standard  solution,  and  /  the  indicator : 

(a)  s^-R:tRS. 

(b)  i-\-R:$jii, 

(c)  S-\-RI:^RS-\-L 

The  recognition  of  the  end-point  depends  upon  the  formation 
of  a  visible  quantity  of  RP ;  this  substance  should  be  incapable 

^  Rlis  not  necessarily  a  compound  of  the  indicator  with  R,  but  even  so  a  definite 
concentration  of  the  free  reagent  R  must  finally  be  present  in  the  solution  in  order 


I04  QUANTITATIVE  CHEMICAL  ANALYSIS 

of  existence  in  the  presence  of  5,  but  after  the  concentration  of 
S  has  been  lowered  to  an  infinitesimal  value,  the  concentration 
of  RI  should  increase  very  rapidly  with  the  further  addition  of 
R.  The  small  amounts  of  RI  formed  locally  in  consequence 
of  imperfect  mixing  should  of  course  react  with  S  and  insure  the 
completion  of  the  main  reaction,  before  the  indicator  becomes 
permanently  affected  by  R.  The  closeness  of  agreement  between 
the  observed  and  the  true  end-point  under  the  conditions  prev- 
alent in  any  specific  case  will  depend  therefore  upon  the  relative 
magnitudes  of  the  equilibrium  constants  of  these  three  reactions. 
In  those  cases,  however,  in  which  the  indicator  combines  with 
the  substance  undergoing  analysis,  the  mechanism  of  the  ti- 
tration may  be  represented  as  follows : 

{a)  S  +/$5/, 

{h)  s-\-Rt:RS, 

(c)  SI+Rt^RS+I. 

Here  the  end-point  is  recognized  by  the  visible  disappearance 
of  SI  from  the  solution.  In  a  few  cases,  in  which  SI  will  not 
react  with  Ry  it  is  necessary  to  use  a  special  procedure ;  as  the 
end-point  is  approached,  the  solution  is  frequently  tested,  a 
drop  at  a  time  on  a  test-plate,  with  a  drop  of  the  indicator  solu- 
tion (see  the  determination  of  ferrous  iron  with  potassium  di- 
chromate  solution).  • 

A  factor  which  is  often  of  the  greatest  importance  is  the  con- 
centration of  the  indicator  in  the  solution.  In  the  reactions  of 
neutralization,  in  which  the  end-point  is  most  frequently  recog- 
nized by  a  sudden  change  from  one  color  to  another  (or  to  color- 
less) in  the  solution,  only  a  very  slight  amount  of  indicator  should 
be  used ;  ^    otherwise  the  entire  quantity  would  not  be  trans- 

to  produce  the  visible  change  which  is  characteristic  of  the  indicator.  In  such 
cases  reaction  (b)  may  be  written  /i  $  I2,  and  the  appearance  of  the  end-point 
is  dependent  upon  the  concentration  of  h,  which  in  turn  depends  upon  that  of  R 
in  the  solution. 

*  In  the  case  of  methyl  orange,  for  example,  the  indicator  is  prepared  by  dis- 
solving 0.02-0.03  g.  of  the  solid  compound  in  100  cc.  of  water,  and  in  any  one 


VOLUMETRIC  ANALYSIS  105 

formed  by  a  slight  enough  excess  of  the  reagent,  and  there  would 
be  a  more  gradual  change  of  color.  On  the  other  hand,  if  RI 
is  a  compound  subject  to  dissociation,  it  is  often  necessary,  in 
order  by  mass  action  to  insure  its  prompt  formation,  to  add  the 
indicator  in  considerable  amount.  (Cf.  the  use  of  starch  in 
iodometry,  and  of  ferric-ion  in  the  titration  of  silver  with 
thiocyanate  solution.) 

Volumetric  vs.  Gravimetric  Methods.  Volumetric  methods 
are  usually  more  rapid  and  convenient  of  execution  than  gravi- 
metric. The  titration  itself  takes  only  a  short  time;  but  the 
preliminary  removal  of  interfering  substances  and  the  trans- 
formation of  the  substance  into  a  form  suitable  for  titration  may 
require  several  hours.  Some  volumetric  processes  are  more 
accurate,  and  others  less  accurate,  than  the  corresponding  gravi- 
metric processes.  They  often  avoid  errors  due  to  solubility, 
to  the  contamination  of  precipitates,  and  to  mechanical  losses, 
which  are  inherent  in  the  gravimetric  processes;  but  they  in- 
volve certain  other  errors  in  the  volumetric  measurement  of 
liquids  and  in  the  determination  of  end-points. 

General  Remarks.  The  student  should  carefully  study  the 
discussion  in  Part  I  concerning  the  sources  of  error  in  the  volu- 
metric measurement  of  liquids.  In  the  way  of  summary,  the 
following  points  may  be  emphasized. 

All  measuring  vessels  should  at  the  same  temperature  be 
based  upon  the  same  standard  of  volume,  and  the  conditions 
observed  in  their  use  should  be  identical  with  those  which  pre- 
vailed in  their  caUbration.  Also,  in  working  with  standard 
solutions,  the  conditions  prevalent  during  standardization  should 
always  be  observed  in  analysis;  this  applies  especially  to  the 
temperature,  to  the  final  volume  after  titration,  and  to  the 
amount  of  indicator  used.  The  value  of  the  solution  is  apt  to 
vary  appreciably  with  the  experimental  conditions. 

titration  only  about  three  drops  of  this  solution  should  be  used.  Counting  20 
drops  to  I  cc,  and  using  3  drops  in  a  titration,  the  amount  of  methyl  orange  actually 
present  is  less  than  0.05  mg. 


io6  QUANTITATIVE  CHEMICAL  ANALYSIS 

It  is  of  course  necessary  to  control  the  value  of  the  standard 
solutions.  They  should  be  preserved  in  closely  stoppered  bottles 
at  a  fairly  constant  temperature,  in  some  instances  in  the  dark. 
Before  use,  the  bottle  should  be  shaken,  to  cause  the  solution  to 
again  take  up  any  liquid  that  may  have  evaporated  and  con- 
densed upon  the  glass;  and,  after  use,  before  replacing  the 
stoppers,  the  necks  and  stoppers  of  the  bottles  should  be  wiped 
dry  with  a  lintless  cloth.  In  addition  to  all  this,  solutions  sub- 
ject to  loss  by  volatilization  or  to  chemical  change,  should  have 
their  value  controlled  at  suitable  intervals  by  restandardization. 

Measuring  vessels  should  be  clean  and  free  from  grease  films, 
and  before  filling,  they  should  be  washed  out  with  three  separate 
small  portions  of  the  solution  to  be  measured,  and  the  washings 
discarded  (through  the  tip,  in  the  case  of  burettes  and  pipettes). 

In  general,  economy  of  time  should  be  practiced  by  calcu- 
lating roughly  the  volume  of  the  standard  solution  which  will 
be  needed  in  a  titration,  and  then  running  in  almost  this  quan- 
tity from  the  burette;  this  can  be  followed  by  a  careful  de- 
termination of  the  end-point.  In  case  such  a  calculation  can- 
not be  made,  it  is  often  worth  while  to  run  a  very  rapid  pre- 
liminary titration,  in  order  to  approximately  fix  this  volume 
(see,  for  example,  the  determination  of  manganese  by  titration 
with  potassium  permanganate). 


VOLUMETRIC  ANALYSIS  107 

A.  NEUTRALIZATION  METHODS 
ACIDIMETRY  AND   ALKALIMETRY 

Fundamental  Principles.  An  aqueous  solution  which  is 
neutral  contains  H+  and  OH"  ions  in  equal  concentration, 
and,  so  long  as  the  solution  is  very  dilute,  the  number  of  mols 
of  water  ionized  in  one  liter  is  o.ooooooi,  or  io~^,  at  25°.  Since 
in  the  ionization  equilibrium,  H2O  :|:  H"**+OH"~,|we  must  have 
(H+)(OH-)  =  ;fe,  it  follows  that  in  dilute  solutions  (H+)(OH-)  = 

If  an  acid  is  added  to  water,  or  to  a  dilute  neutral  solution, 
in  sufficient  amount  to  increase  the  H+-ion  concentration  from 
io~^  to  iQ-^,  then  the  OH^-ion  concentration  will  fall  to  io~^. 
Only  0.0000009  mol  of  H+  ion  would  be  required  to  produce 
such  a  change  in  a  liter  of  neutral  solution,  and  this  amount  is 
contained  in  about  o.oi  cc.  of  tenth-normal  hydrochloric  acid, 
while  the  addition  of  only  o.i  cc.  to  one  liter  would  increase  the 
H+-ion  concentration  from  io~^  to  io~^. 

The  indicators  commonly  used  to  determine  the  end-point 
in  processes  of  acidmetry  and  alkalimetry  are  rather  compli- 
cated organic  compounds,  each  capable  of  existence  in  two 
forms  of  different  color,  —  these  forms  being  mutually  con- 
vertible, one  into  the  other,  at  specific  H+-  and  OH~-ion  con- 
centrations. 

The  color  changes  which  are  characteristic  of  these  indicators 
do  not  as  a  rule  occur  in  exactly  neutral  solutions.  The  accom- 
panying table  gives  for  some  of  the  more  common  indicators, 
the  H+-ion  concentrations  at  which  the  color  changes  occur. 

It  might  seem,  at  first  thought,  that  only  an  indicator  which 
changes  exactly  at  the  neutral  point  would  be  suitable,  but  this 
is  by  no  means  true.  In  titrating  a  strong  acid  against  a  strong 
base  a  few  hundredths  of  a  cubic  centimeter  of  tenth-normal 
acid  or  alkali  in  excess  will  carry  the  concentration  of  hydrogen 
or  hydroxide  ion  so  far  to  one  side  of  the  neutral  point  that  any 


io8  QUANTITATIVE  CHEMICAL  ANALYSIS 

of  the  indicators  for  which  the  characteristic  point  lies  between 
io~^  and  io~^  will  give  a  sharp  and  accurate  end-point. 

In  titrating  a  weak  acid  with  a  strong  base,  e.g.  acetic  acid 
with  sodium  hydroxide,  the  acetate  ions  from  the  highly  ionized 
sodium  acetate  drive  the  reaction,  HC2H302:$H++C2H302~, 
far  to  the  left,  long  before  an  equivalent  quantity  of  the  bas3 
has  been  added;  the  concentration  of  hydrogen  ion  becomes 
exceedingly  low  (<io"^  but  >io-^)  before  the  major  portion  of 
the  acetic  acid  has  had  its  hydrogen  replaced.  In  such  a  case 
the  change  in  color  for  Methyl  orange  (io"^io"^)  will  appear 
gradually,  during  the  addition  of  a  cubic  centimeter  or  more  of 
the  alkali,  long  before  an  equivalent  quantity  has  been  run  in, 
and  no  sharp  end-point  will  be  indicated.  If  phenolphthalein 
(io~*-io~^)  is  used,  however,  the  change  in  color  will  not  occur 
until  the  true  neutral  point,  (H+)  =  lo"^,  has  been  passed ;  a  very 
slight  excess  of  alkali  will  then  reduce  the  concentration  of  the 
hydrogen  ion  far  below  lo"^,  a  sharp  end-point  will  be  indicated, 
and  the  total  alkali  added  will  correspond  very  exactly  to  that 
required  for  the  actual  neutralization  of  the  acid. 

When  a  weak  base  is  titrated  with  a  strong  acid,  e.g.  ammonium 
hydroxide  with  hydrochloric  acid,  the  conditions  are  reversed 
and  such  indicators  as  methyl  orange,  methyl  red,  or  cochineal, 
which  change  color  in  a  faintly  acid  solution  {i.e.  H+  >  io~0  are 
most  suitable. 

The  following  rules  should  be  observed  in  the  use  of  neutraliza- 
tion methods  for  the  determination  of  acids  and  bases : 

1.  In  the  titration  of  a  strong  acid  with  a  strong  base,  or  vice 
versa,  use  any  indicator  in  the  list,  from  methyl  orange  to  phenol- 
phthalein. 

2.  In  the  titration  of  any  acid  other  than  a  strong  mineral 
acid  with  a  strong  base,  use  phenolphthalein,  trinitrobenzene, 
or  a  similar  indicator. 

3.  In  the  titration  of  a  weak  base  with  a  strong  acid,  use 
methyl  orange,  Congo  red,  or  a  similar  indicator. 

4.  Do  not  attempt  to  titrate  a  weak  base  against  a  weak  acid. 


VOLUMETRIC  ANALYSIS 


109 


\ 

+ 

+ 

I 

s 

'0 

+ 

+ 

+ 

1 

+ 

M 

+ 

+ 

+ 

+ 

+ 

1 

+ 

+ 

+ 

+ 

+ 

'+ 

+ 

+ 

+ 

+ 

+ 

+ 

S 

+ 

+ 

> 

I 

+ 

+ 

. 

+ 

+ 

+ 

+ 

I 

+ 

+ 

1 

I 

+ 

+ 

+ 

+ 

+ 

+ 

1 

1 

+ 

+ 

+ 

"k 

+ 

+ 

+ 

+ 

Sfii 

+ 

> 

1 

+ 

+ 

+ 

J 

I 

+ 

+ 

i 

>> 

6 

1 

+ 

+ 

+ 

1 

2. 

a 

\ 

+ 

1 

+ 

1 

&4 

^> 

1^ 

2 

+ 

+ 

a 

\ 

St 

0 

+ 

+ 

> 

> 

+ 

+ 

1 

\ 

m 

+ 

+ 

s 

n 

+ 

+ 

1 

\ 

+ 

+ 

+ 

+ 

+ 

+ 

+ 

I 

+ 

+ 

+ 

- 

i2 

^1 

n 

IS 

+ 

li 

1 

0 

1 

it 
1 

i 
1 

2 

5 

1 

M 

1 

1 

1 

i 

3 

d 
.a 

1 
1 

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i 
1 

no  QUANTITATIVE  CHEMICAL  ANALYSIS 

The  sensitivity  of  a  given  indicator  may  vary  widely  with 
the  temperature  or  dilution,  and  for  that  reason  the  end-point 
should  be  determined  in  about  equal  volumes  of  solution,  both 
in  standardization  and  in  analysis  ;i  and  when,  as  is  often  neces- 
sary, a  solution  is  to  be  titrated  at  the  boiHng  temperature,  the 
standardization  also  should  be  carried  out  at  that  temperature. 
And,  since  a  slight  amount  of  acid  or  of  alkali  is  necessary  to 
transform  the  indicator,  the  constancy  of  this  amount  should  be 
insured  by  always  using  a  fixed  quantity  of  the  indicator  solu- 
tion, say  3  drops. 

Methyl  orange  solution  is  most  readily  prepared  by  dissolv- 
ing 0.02-0.05  g-  of  the  solid  compound  (also  known  as  Orange  III) 
in  a  very  little  alcohol  and  diluting  with  water  to  100  cc.  It 
is  often  used  in  the  titration  of  strong  acids  and  bases,  and 
especially  in  the  determination  of  weak  bases,  as  ammonium 
hydroxide.  It  is  also  serviceable  in  titrating  with  a  strong  acid 
the  salts  of  very  feeble  acids,  as  carbonates,  sulphides,  borates, 
etc. ;  in  such  cases  the  acids  set  free  are  too  slightly  ionized  to 
affect  the  indicator,  and  the  change  to  pink^does  not  occur  until 
the  standard  acid  has  been  added  in  very  slight  excess.  In  reahty, 
the  strong  acid  simply  removes  the  ©H^-ions  which  accumu- 
late in  the  solution  in  consequence  of  the  hydrolysis  of  the  salt, 
and  thus  drives  the  hydrolysis  to  completion :  e.g. 

Na2Si03+2  H0H:^2  Na++H2Si03+2  OH- 

The  use  of  methyl  orange  should  be  confined  to  cold  solu- 
tions, which  as  a  rule  are  to  be  titrated  by  standard  acid  or  alkali 
of  about  normal  or  half-normal  concentration. 

Phenolphthalein  solution  is  prepared  by  dissolving  about 
I  g.  of  the  soHd  in  55-60  cc.  of  95%  alcohol,  and  diluting  the 
solution  to  100  cc.  This  indicator  is  especially  adapted  to  the 
titration  of  sHghtly  ionized  acids,  as  acetic  and  tartaric  acids. 
(It  should  not  be  used  in  connection  with  weak  bases,  such  as 

1  In  this  connection,  see  G.  P.  Baxter,  Jour.  Amer.  Chem.  Soc,  Vol.  j(5,  p.  656 
(1914). 


VOLUMETRIC  ANALYSIS  iii 

ammonia.)  It  is  decolorized  even  by  carbonic  acid,  which  there- 
fore must  be  removed  by  heating  when  other  substances  are 
being  determined ;  unUke  methyl  orange,  it  is  sensitive  in  boil- 
ing hot  solutions. 

Standard  Aci^  Solutions.  These  are  generally  prepared  from 
hydrochloric  or  sulphuric  acid.  Hydrochloric  acid  is  preferable 
to  sulphuric  in  the  titration  of  barium  hydroxide ;  but,  on  the 
other  hand,  solutions  of  sulphuric  acid  can  be  used  in  titrations 
at  the  boiling  temperature  without  danger  of  loss.  Both  acids 
may  be  used  with  any  indicator  in  the  table  from  methyl  orange 
to  phenolphthalein. 

Standard  Alkali  Solutions.  These  are  mostly  prepared  from 
sodium  hydroxide,  but  sodium  carbonate  and  barium  hydroxide 
are  sometimes  used.  Sodium  hydroxide,  if  free  from  carbonate, 
may  be  used  with  any  indicator ;  but  its  solutions  absorb  car- 
bon dioxide  from  the  air.  Pure  sodium  carbonate,  accurately 
weighed  out,  may  be  dissolved  in  water  and  diluted  to  give  a 
standard  solution,  but  its  usefulness  may  be  impaired  by  the 
liberation  of  carbonic  acid.  Barium  hydroxide  solutions  are 
free  from  carbonate,  since  any  carbon  dioxide  absorbed  is  pre- 
cipitated as  insoluble  barium  carbonate;  this,  however,  lowers 
the  concentration  of  the  base.  Barium  hydroxide  may  be  used 
with  any  indicator,  from  methyl  orange  to  phenolphthalein, 
but  it  is  not  soluble  much  above  one  fourth  normal ;  if  a  more 
concentrated  solution  of  carbonate-free  alkali  is  desired,  how- 
ever, a  sodium  hydroxide  solution  may  be  freed  from  carbonate 
by  means  of  a  slight  excess  of  barium  chloride,  and  then  stand- 
ardized. Carbonate-free  solutions  should  be  protected  from 
carbon  dioxide  by  means  of  a  soda-lime  absorption  tube. 

THE  PREPARATION  AND  STANDARDIZATION  OF  APPROXI- 
MATELY HALF-NORMAL  SOLUTIONS  OF  HYDROCHLORIC 
ACID  AND   SODIUM  HYDROXIDE 

Procedure.  Pour  into  a  loo  cc.  measuring  cyHnder  a  volume 
of  hydrochloric  acid  {e.g.  82.5  cc.  of  the  acid  of  sp.  gr.  1.19) 


112  QUANTITATIVE  CHEMICAL  ANALYSIS 

suflSicient  to  contain  36.5  g.  of  hydrogen  chloride,  transfer  it 
quantitatively  to  a  one-liter  measuring  flask,  and  dilute  to  the 
mark  with  distilled  water.  Pour  this  solution  without  loss  into 
a  clean,  well-drained  2.5  liter  bottle,  refill  the  measuring  flask 
with  water,  and  pour  this  also  into  the  bottle.  Finally  mix  the 
solution  thoroughly  by  vigorous  shaking  in  the  stoppered  bottle. 

In  a  large  beaker,  dissolve  in  water  about  42  g.  of  stick  sodium 
hydroxide,  weighed  on  a  rough  balance,  dilute  this  solution  also 
to  2  liters,  and  shake  the  mixture  in  the  stoppered  bottle. 

Allow  these  solutions  to  stand,  best  overnight,  until  they  have 
attained  the  temperature  of  the  laboratory. 

Fill  two  clean  burettes  with  the  respective  solutions,  first 

thoroughly  rinsing  each  burette  with  one  10  cc,  and  two  5  cc. 

portions  of  the  corresponding  solution,  each  time  draining  off 

the  wash  liquid  through  the  tip.     See  that  the  tips  are  free  from   * 

air,  and  then  record  the  exact  reading  of  each  burette.     Run 

out  about  20  cc.  of  the  acid  into  a  500  cc.  Erlenmeyer  flask,  add 

60  cc.  of  water  and  3  drops  of  methyl  orange  solution,  and  then 

from  the  other  burette  run  in  alkali,  with  shaking,  until  the  color 

of  the  solution  changes  from  pink  to  yellow.     Wash  down  the 

inside  of  the  flask  with  a  little  water,  replace  it  under  the  first 

burette,  and  add  acid  until  the  solution  turns  pink.     In  this  way, 

accurately  adjust  the  mixture  until  a  drop  of  either  acid  or 

alkali  vn\\  cause  a  definite  change  of  color.     Select  the  point  at 

which  the  faintest  tinge  of  pink  appears  as  the  end-point  of  the 

titration;    later  on,  always  titrate  to  this  point.     Allow  about 

30  seconds  for  afterflow,  and  then  record  the  readings  of  the 

burettes.     Having  corrected  the  readings  in  accordance  with 

the  burette  cahbrations,  and  with  the  temperature  if  necessary 

(see  Part  I),  calculate  the  ratio  of  the  solutions  as  shown  in  the 

foiling  example : 

cc.  acid       21.53  r      .J  r    n    T 

— -:= — ^^=1.022  cc.  of  acid  per  i.cx^o  cc.  of  alkah. 

cc.  alkah     21.07 

Fill  the  burettes  again,  and  repeat  the  operation.  The  dupli- 
cate ratios  should  check  within  two  parts  in  one  thousand. 


VOLUMETRIC  ANALYSIS  113 

With  this  ratio  accurately  established,  standardize  the  hydro- 
chloric acid  solution  as  follows : 

{a)  By  Titration  against  Pure  Sodium  Carbonate.  If  a  suit- 
able oven  is  available,  dry  the  salt  on  a  watch  glass  for  an  hour 
at  130-150°;  otherwise  heat  about  10  g.  of  pure  sodium  car- 
bonate in  a  small  porcelain  dish,  on  a  wire  gauze  over  a  small 
Bunsen  flame,  for  one  half  hour,  and  then  allow  the  salt  to  cool 
in  a  desiccator.  Transfer  the  cold  salt  to  a  dry,  well-stoppered 
weighing  tube,  and  weigh  out  into  500  cc.  Erlenmeyer  flasks 
two  portion^of  0.5-0.6  g.  each,^  l^cording  the  exact  weights. 
Add  to  each  sample  80  cc.  of  water  and  3  drops  of  methyl  orange 
solution,  and  shake  until  dissolved.  With  both  burettes  full 
and  the  readings  recorded,  run  in  the  acid  with  shaking,  until 
the  solution  just  turns  pink.  Wash  down  the  inside  J^f  the 
flask,  and,  if  the  solution  turns  yellow,  add  acid  a  drop  at  a 
time,  until  the  faintest  tinge  of  pink  returns.  (If  too  much 
acid  is  added,  the  excess  can  be  determined  by  means  of  the 
alkali  in  the  other  burette,  since  the  ratio  of  the  solutions  is 
known.)  After  a  few  moments,  read  the  burettes  and  record 
the  readings.  From  the  data  obtained  in  each  case,  calculate 
the  normality  factor  of  the  acid;  and  also  that  of  the  alkali, 
using  the  mean  of  the  dupKcate  values  obtained  in  determining 
the  ratio  of  the  two  solutions.  The  two  results  should  agree 
within  two  parts  in  one  thousand.  ^ 

{h)  Gravimetrically  with  Silver  Nitrate.  Measure  out  accu- 
rately from  a  pipette  10.00  cc.  of  the  acid  into  each  of  two 
300-cc.  beakers,  and  dilute  in  each  case  with  150  cc.  of  waty. 
Precipitate  the  chlorine  from  these  solutions  with  silver  nitrate 
according  to  th^procedure  given  in  Part  II,  and  filter  the  silver 
chloride  off  through  Goofc  crucibles,  prepared  and  weighed  as 
there  indicated.  Wash  the  precipitates  with  hot  water ||mi til 
free  from  soluble  silver  salts,  dry  at  120-130°  to  constant  \<^^ht, 
and  from  the  weight  of  silver  chloride  found  in  each  case  calcu- 

*  The  weights  of  samples  in  this  book  are  based  upon  the  use  of  30-cc.  burettflfe. 
If  so-cc.  burettes  are  used,  it  is  better  to  take  samples  f  as  large. 


114  QUANTITATIVE  CHEMICAL  ANALYSIS 

late  the  normality  factor  of  the  acid.  The  duplicate  values 
should  agree  very  closely  with  one  another,  and  also  with  those 
previously  found  by  titration  against  sodium  carbonate. 

Notes.  —  i.  Although  silver  chloride  is  insoluble,  the  normality  factor 
of  the  hydrochloric  acid  may  nevertheless  be  calculated  directly  from  the 
weight  of  the  precipitate  obtained.  For  example,  if  lo.oo^cc.  of  the  acid 
were  found  to  yield  0.7317  g.  of  AgCl,  then  (since  an  equal  volume  of  normal 
HCl  would  yield  1.4334  g.  of  AgCl)  the  normaUty  factor  of  the  acid  is 
0.7317/1.4334,  or  0.5105. 

In  the  same  way,  the  normality  factor  of  the  acid  may  be  calculated  in 
the  case  of  Method  {a).  For  Sample,  if  0.5682  g.  of  Na2C03  require 
21.00  cc.  of  acid,  the  normality  factor  is  equal  to  0.5682/1.1130=0.5105 
(1.1130  g.  is  the  amount  of  Na2C03  contained  in  21.00  cc.  of  the  normal 
solution). 

If  jj^has  previously  been  found,  for  example,  that  i.ooo  cc.  of  alkali 
solution  is  equivalent  to  1.022  cc.  of  the  acid,  then  it  follows  that  the  nor- 
mality factor  of  the  alkali  is  0.5105X1.022  =  0.5217. 

2.  If  it  is  desired  to  prepare  solutions  of  exactly  one  half  normal  con- 
centration, slightly  stronger  solutions  are  first  prepared,  and,  after  stand- 
ardization, they  are  diluted  with  the  calculated  volume  of  water.  For 
example,  the  0.5105  N  acid  should  be  diluted  according  to  the  proportion, 
0.5105  : 0.5000=01; :  1000,  and  the  0.5217  iV  alkali  according  to  the  proportion, 
0.5217 : 0.5000= y:  1000;  i.e.  one  liter  of  each  solution  should  have  added 
to  it  21.0  cc.  and  43.2  cc.  of  water,  respectively.  The  water  is  added  from 
a  burette.  After  dilution  the  solutions  should  be  thoroughly  shaken,  and 
then  restandardized. 

3.  In  fixi^^  the  end-point  of  the  titration,  it  is  somewhat  easier  to  recog- 
nize the  change  from  yellow  to  pink  than  that  from  pink  to  yellow.  Which- 
ever change  is  selected  as  the  end-point,  however,  should  be  consistently 
made  use  of  in  all  subsequent  titrations. 

*  4.  There  has  been  much  discussion  concerning  the  compound  which  is 
best  suited  for  use  in  the  standardization  of  acids.  Although  calcium  car- 
bonate has  many  advocates  (see  Part  V,  Problem  54),  it  is  more  customary 
to  use  sodium  carbonate.  This  salt  has  thikadvantage  of  being  soluble  in 
water,  and  it  can  now  be  purchased  sufficiently  pure  for  the  purpose.  The 
puijpalt  can  easily  be  prepared  by  heating  the  recrystallized  bicarbonate 
at  280-300°,  when  it  decomposes  according  to  the  equation, 

2  NaHC03=Na2C03+H20-hC02. 
^  5.  Of  course  it  is  possible  to  standardize  the  alkali  solution  directly, 
rather  than  the  acid.    With  phenolphthalein  as  the  indicator,  this  may  be 


VOLUMETRIC  ANALYSIS  115 

done  against  any  of  the  following  pure  acids :  oxalic  acid,  H2C2O4 .  2  H2O ; 
acid  potassium  oxalate,  HKC2O4 .  H2O;  potassium  tetroxalate,  H3K(C204)2 
.  2  H2O ;  succinic  acid,  H2C4H4O4 ;  acid  potassium  tartrate,  HKC4H4O6 ; 
acid  potassium  phthalate,  HKC8H4O4.  The  last  two  are  probably  the  most 
suitable,  since  they  are  monobasic  acids  of  high  molecular  weight,  do  not 
contain  water  of  crystallization,  and  are  easily  obtainable  in  a  pure  condition. 

6.  While  it  is  permissible  to  standardize  hydrochloric  acid  solutions 
(providing  they  are  free  from  interfering  impurities)  gravimetrically  with 
silver  nitrate,  solutions  of  sulphuric  acid  should  not  be  standardized  by 
precipitation  as  barium  sulphate ;  in  this  case  the  purity  of  the  precipitate 
is  too  open  to  question.  ^ 

7.  Before  beginning  work  in  volumetric  analysis,  the  student  should 
read  Section  VI  of  Part  I,  as  well  as  the  general  discussion  at  the  beginning 
of  Part  III. 


ii6  QUANTITATIVE  CHEMICAL  ANALYSIS 

THE  DETERMINATION  OF  THE  ALKALINE  VALUE  OF  SODA 

The  sample  may  be  one  of  commercial  soda,  or  it  may  be  an 
artificial  mixture  of  sodium  carbonate  and  sodium  chloride. 

Procedure.  Weigh  out  roughly  on  a  clean  watch  glass  about 
5  g.  of  the  soda,  dry  it  for  i  hour  at  iio°,  and  allow  to  cool  in  a 
desiccator.  Now  accurately  weigh  the  sample  on  the  watch  glass, 
transfer  it  quantitatively  to  a  beaker,  and  dry  and  weigh  the 
watch-glass  in  order  to  arrive  at  the  exact  weight  of  the  sample. 

Warm  the  sample  with  foo  cc.  of  water  and,  after  cooling, 
filter  from  any  insoluble  matter,  thoroughly  washing  the  filter 
with  cold  water  and  receiving  the  filtrate  and  washings  in  a 
clean  250  cc.  measuring  flask.  At  the  room  temperature,  dilute 
to  the  graduation  mark  with  water  (see  p.  45,  Parallax),  adding 
the  latter  gradually  and  with  shaking ;  and  then,  using  a  clean 
dry  beaker,  thoroughly  mix  the  solution  by  pouring  it  back  and 
forth  from  one  vessel  into  the  other. 

Measure  out  two  or  three  25  cc.  portions  of  the  liquid,  by 
means  of  a  pipette  previously  rinsed  with  the  solution,  into  500 
cc.  Erlenmeyer  flasks,  touching  the  tip  of  the  pipette  for  final 
drainage  each  time  on  the  inner  wet  surface  of  the  vessel  (un- 
less it  was  calibrated  otherwise).  Add  50  cc.  of  water  and  3 
drops  of  methyl  orange  solution,  and  titrate  with  the  standard 
acid,  as  in  the  standardization  of  the  latter. 

From  the  (corrected)  volumes  of  acid  and  alkali  used,  re- 
membering that  only  one  tenth  of  the  sample  was  used  in  each 
titration,  calculate  the  quantity  of  alkali  present  in  terms  of  pure 
sodium  carbonate,  and  report  the  percentage  of  the  latter  found. 

Notes.  —  i.  Let  us  assume,  for  example,  that  5.890  g.  of  soda  were 
used  in  the  preparation  of  250.0  cc.  of  solution,  and  that  25.75  cc.  of  0.5105  N 
aci<|[^d  4.13  cc.  of  0.5217  N  alkali  were  used  in  the  titration  of  25.00  cc.  of 
this  solution.  Then  it  follows  that  5.890  g.  of  the  soda  are  equivalent 
to  ioX(25. 75X0.5105— 4.i3Xo.52i7)  =  i09.90  cc.  of  N  acid;  and,  since 
this  volume  of  normal  acid  would  neutralize  an  equal  volume  of  normal 
alkali,  therefore  the  5.890  g.  of  soda  contained  0.053X109.90=5.8246  g., 
or  98.9%  of  Na2C03. 


VOLUMETRIC  ANALYSIS  117 

2.  Soda  is  apt  to  contain  varying  amounts  of  water,  and,  for  the  sake 
of  uniformity,  the  sample  to  be  analyzed  is  ordinarily  dried  for  a  specified 
time  at  110° ;  though  this  treatment  is  insufl&cient  to  completely  remove  the 
moisture. 

3.  Crude  Solvay  soda  usually  contains  as  impurities,  besides  a  small 
amount  of  sandy  grit,  also  small  quantities  of  sodium  chloride  and  sulphate, 
and  either  the  bicarbonate  or  hydroxide.  Either  of  the  latter  two  has  an 
alkaline  value  of  its  own,  but  nevertheless  the  alkaline  value  is  ordinarily 
expressed  in  terms  of  sodium  carbonate  alone. 

4.  The  chemical  action  in  the  titration  of  sodium  carbonate  with  hydro- 
chloric acid  consists  essentially  of  two  stages;  first  the  neutralization  of 
hydroxide-ion,  which  accumulates  in  the  solution  in  consequence  of  hydroly- 
sis, and  second  the  conversion  of  the  bicarbonate-ion  of  the  first  reaction 
into  non-ionized  carbonic  acid.  The  carbonic  acid  itself  is  too  feeble  to 
have  any  effect  upon  methyl  orange ;  it  decomposes  mostly  into  water  and 
carbon  dioxide. 

(a)  NaaCOa+HOH  $  2  Na++HC03-+0H-  \  ^  h  O  • 

HCl  :§  CI-      -h   H+   /        '  _' 

{b)  NaHCOa  $  Na+-|-HC03l\  _  w  m 

HCi:§:Cl-H-H-       1|-H,C03. 

5.  For  other  methods  of  analyzing  soda,  the  student  is  referred  to 
Part  V,  Problems  60,  61,  and  96. 


li8  QUANTITATIVE  CHEMICAL  ANALYSIS 


THE  DETERMINATION   OF  AVAILABLE   HYDROGEN-ION  IN 

AN  ACID 

The  sample  may  be  oxalic  acid,  acid  potassium  oxalate,  po- 
tassium tetroxalate,  succinic  acid,  acid  potassium  phthalate, 
potassium  bitartrate,  or  something  similar. 

Procedure.  Weigh  out  accurately  into  500  cc.  Erlenmeyer 
flasks  0.6-0.7  g-  portions  of  the  unknown  acid,  and  treat  each  as 
follows.  Warm  the  sample  with  80  cc.  of  water,  add  3  drops  of 
phenolphthalein  solution,  and  with  shaking,  run  in  the  stand- 
ard alkali  until  a  pink  color  persists.  Cautiously  add  the  stand- 
ard acid,  from  a  second  burette,  until  the  color  just  vanishes, 
and  then  an  excess  of  0.3-^.4  cc.  Heat  the  solution  for  3  minutes 
at  the  boiling  temperature,  and,  if  the  color  reappears  add  acid 
as  before  to  the  hot  solution,  again  0.3-0.4  cc.  in  excess,  and  boil 
again  for  3  minutes.  Continue  this  treatment  until  the  pink 
color  is  not  restored  by  boiUng  for  3  or  4  minutes.  Finally  add 
alkali  until  the  faintest  tinge  of  pink  returns,  then  2  drops  of 
acid,  and  boil,  and  if  no  color  returns,  finish  the  titration  with 
alkali  at  the  boiling  temperature. 

From  the  data  obtained,  calculate  the  normality  factor  of  a 
solution  containing  i.ooo  g.  of  the  unknown  acid  in  100  cc. 
The  results  should  check  within  two  parts  in  one  thousand. 

Notes.  —  i.  Although,  theoretically  speaking,  the  same  indicator  (and 
temperature)  should  be  used  in  standardization  and  analysis,  the  error  in- 
volved in  the  non-observance  of  this  principle  is  in  this  instance  practically 
negligible. 

2.  Since  phenolphthalein  is  sensitive  to  carbonic  acid,  and  since  the 
ordinary  standard  alkali  always  contains  carbonate,  it  is  necessary  to  expel 
this  acid  by  boiling  the  solution.  In  a  cold  dilute  solution  of  sodium  bi- 
carbonate which  is  saturated  with  carbon  dioxide,  phenolphthalein  is  color- 
less ;  i.e.  in  cold  dilute  solutions  of  sodium  carbonate  this  indicator  becomes 
colorless  as  soon  as  the  carbonate  has  been  transformed  into  bicarbonate  by 
the  addition  of  acid  (see  Note  4  of  the  preceding  exercise).  When  this 
solution  is  heated,  the  bicarbonate  is  partially  hydrolyzed, 

Na+HC03-+H+0H-  t>  Na+-f-0H--i-H2C0a, 


VOLUMETRIC  ANALYSIS  119 

and  the  solution  loses  carbon  dioxide  and  becomes  alkaline.  This  behavior 
will  continue  durmg  the  titration,  as  long  as  there  is  any  bicarbonate  left 
in  the  solution. 

3.  If  present  at  any  appreciable  concentration,  hydrogen  chloride  is 
more  or  less  volatilized  from  boiling  aqueous  solutions ;  for  this  reason  the 
acid  should  be  added  each  time  in  very  slight  excess. 

4.  When  a  large  number  of  determinations  are  to  be  made  with  phenol- 
phthalein  as  the  indicator,  it  is  well  worth  while  to  prepare  and  standardize  a 
carbonate-free  alkali  solution.  The  acid  solutions  in  such  cases  should  be 
boiled  (unless  there  is  danger  of  loss),  to  free  them  from  carbonic  acid,  or 
they  may  be  made  up  with  freshly  boiled  water,  and  titrated  hot  with  car- 
bonate-free alkali.  For  the  greatest  accuracy,  a  specially  constructed,  steam 
jacketed  titration  vessel  may  be  employed,  in  which  case  carbonate-free 
alkali  is  not  really  needed. 


I20  QUANTITATIVE  CHEMICAL  ANALYSIS 

THE    DETERMINATION    OF    PROTEIN    NITROGEN    BY  THE 
KJELDAHL  METHOD 

Principle.  When  an  organic  substance  is  heated  with  con- 
centrated sulphuric  acid,  especially  in  the  presence  of  an  oxygen 
carrier,  the  organic  substance  is  completely  decomposed,  and 
any  protein  (or  other  similarly  combined)  nitrogen  is  converted 
into  ammonia.  This  at  once  combines  with  acid  to  form  am- 
monium acid  sulphate,  NH4HSO4,  which  remains  in  solution 
in  the  sulphuric  acid.  Upon  diluting  the  mixture  with  water 
and  adding  sodium  hydroxide  in  excess,  the  ammonia  is  Hberated, 
and  can  be  distilled  over  and  collected  in  a  known  volume  of 
standard  acid,  which  it  partially  neutralizes.  By  titrating  the 
excess  of  acid  with  a  standard  alkali,  the  volume  of  the  standard 
acid  neutralized  by  the  ammonia  can  be  found,  and  from  the 
data  obtained  the  percentage  of  nitrogen  in  the  sample  may  be 
calculated. 

Procedure.  Accurately  weigh  out  from  a  weighing  tube, 
upon  separate  sheets  of  quantitative  filter  paper,  two  samples 
of  about  I  g.  each  of  the  substance  to  be  analyzed.  Wrap  each 
sample  carefully  in  the  paper,  and  introduce  the  bundle  into  a 
clean  500  cc.  Kjeldahl  flask.  To  each  flask  add  about  0.5  g.  of 
powdered  copper  sulphate,  and  25  cc.  of  concentrated  sulphuric 
acid.  See  tljtat  the  samples  are  thoroughly  wet  by  the  acid,  and 
then  place  the  flasks  on  the  digestion  rack  in  the  Nitrogen  Lab- 
oratory, with  the  necks  resting  in  the  circular  openings  of  the 
lead  ventilating  pipe;  place  the  flasks  in  unoccupied  positions 
as  near  as  possible  to  one  of  the  exhaust  flues.  Heat  gently 
until  frothing  ceases,  add  10  g.  (weighed  roughly)  of  potassium 
sulphate,  or  an  equivalent  weight  of  sodium  sulphate,  and 
heat  to  gentle  ebuUition  for  two  or  three  hours  until  the  Kquid 
is  of  a  clear  green  color,  without  any  trace  of  brown  (do  not 
allow  the  flame  to  reach  above  the  surface  of  the  liquid). 
Continue  the  heating  for  half  an  hour  longer,  and  allow 
to  cool. 


VOLUMETRIC  ANALYSIS  121 

While  the  flasks  are  cooling,  accurately  measure  from  a  burette 
two  30.00-cc.  portions  of  0.5  N.  hydrochloric  acid,  into  400-cc. 
Erlenmeyer  flasks,  and  add  to  each  about  25  cc.  of  distilled  water. 
Place  these  flasks  under  the  distilhng  apparatus,  so  that  the 
delivery  tubes  just  dip  into  the  acid  solutions. 

After  cooling,  carefully  dilute  the  contents  of  the  digestion 
flasks  with  150  cc.  of  distilled  water,  and  cool  again.  Carefully 
pour  down  the  inclined  neck  of  each  flask,  so  that  it  shall  not 
mix  with  the  acid  solution,  75  cc.  of  sodium  hydroxide  solution 
(300  g.  of  NaOH  per  liter).  Place  the  flasks  on  the  distilling 
rack,  add  one  or  two  pieces  of  granulated  zinc,  and  quickly 
connect  the  flasks  with  the  distilling  heads,  using  well-fitting 
rubber  stoppers.  Finally,  mix  the  contents  of  each  flask  by 
gently  rotating  it,  and  then  begin  to  heat  the  mixture. 

Distill  off  about  two  thirds  of  the  contents  of  each  flask,  with 
great  care  that  they  do  not  boil  over.  The  distillation  will  re- 
quire about  45  minutes.  Disconnect  the  distilling  flasks  and 
rinse  out  the  delivery  tubes  into  the  receiving  flasks  with  a 
little  distilled  water.  Add  3  drops  of  methyl  orange  to  each  of 
the  receiving  flasks,  and  titrate  the  contents  with  0.5  N.  sodium 
hydroxide. 

From  the  data  obtained,  calculate  the  percentage  of  nitrogen 
in  the  sample. 

Notes.  —  i.  The  sulphuric  acid  hydrolyzes  the  NH2-group,  to  give 
ammonia,  and  also  acts  as  an  oxidizing  agent,  converting  the  organic  matter 
into  carbon  dioxide,  water,  or  other  volatile  products.     For  example, 

CO(NH2)2+H204-2  H2S04=C02+2  NH4HSO4; 

C6Hio06+wH2S04=6  C-hs  H2O+W  H2SO4; 
and  C+2  H2S04=C02+2  H2O4-2  SO2. 

2.  The  CUSO4  gives  up  oxygen  more  readily  to  the  organic  matter  than 
the  H2SO4  does ;  but  the  H2SO4  then  reoxidizes  the  copper  so  that  at  the 
end  of  the  operation  the  copper  is  still  present  as  copper  sulphate.  That 
is  to  say,  the  copper  salt  acts  catalytically  as  an  oxygen  carrier. 

3.  Mercuric  sulphate  is  often  used  instead  of  copper  sulphate  as  an 
oxygen  carrier,  a  few  small  globules  of  metallic  mercury  being  added  to 


122  QUANTITATIVE  CHEMICAL  ANALYSIS 

the  acid  digestion  mixture.  Although  the  mercury  salt  is  somewhat  more 
efficient,  it  tenaciously  retains  ammonia,    as 

HaN-Hg-O-SOa-O-Hg-NHa, 

even  in  the  presence  of  an  excess  of  hot  alkali,  and  it  is  therefore  necessary 
to  add  also  a  large  excess  of  sodium  sulphide.  This  converts  all  the 
mercury  into  HgS,  which  combines  with  Na2S  to  form  soluble  Hg(SNa)2 
and  the  ammonia  is  liberated. 

4.  The  K2SO4  forms  with  the  acid  KHSO4,  and  this  serves  to  raise  the 
boiling  point  of  the  sulphuric  acid ;  the  higher  temperature  hastens  the  di- 
gestion. 

5.  The  flask  is  provided  with  a  long  neck  in  order  that  the  acid  fumes, 
which  would  otherwise  be  lost,  may  condense  and  run  back  into  the  digestion 
mixture. 

6.  After  the  acid  solution  has  been  diluted,  it  is  specifically  lighter  than 
the  NaOH  solution  used ;  upon  pouring  the  latter  carefully  down  the  neck 
of  the  incUned  flask,  it  sinks  to  the  bottom  and  leaves  the  surface  of  the 
liquid  still  acid,  thus  preventing  the  loss  of  ammonia  at  this  stage  of  the 
procedure.  The  contents  should  not  be  mixed  until  after  the  flask  has 
been  tightly  connected  with  the  distilling  head. 

7.  The  granulated  zinc  is  added  in  order  to  prevent  bumping  during 
the  distillation ;  the  zinc  dissolves  slowly  in  the  alkaline  solution,  with  the 
evolution  of  hydrogen.  Fragments  of  pumice  stone  or  of  platinum  are 
often  used  for  the  same  purpose. 

8.  If  nitrates  are  present  in  the  sample  {e.g.  a  fertilizer),  and  it  is  desired 
to  determine  the  total  nitrogen,  the  procedure  may  be  modified  as  follows : 
Thoroughly  wet  the  sample  in  the  flask  with  25  cc.  of  concentrated  sul- 
phuric acid,  in  which  one  gram  of  salicylic  acid,  f  C6H4<^  COOh)  ^^^  Previ- 
ously been  dissolved ;  this  reacts  with  the  nitric  acid  to  form  nitrosaHcylic 

/NO2 
acid,  CeHs — OH     .     Next  add  slowly,  with  frequent  shaking,  10  g.  of 
\COOH 

powdered  sodium  thiosulphate,  which  reduces  the  nitrosalicylic  acid  to 

/NH2 
aminosalicyhc  acid,  CeHs — OH    .    Now  add  0.5  g.  of  powdered  copper 

\COOH 
sulphate    and    complete    the    determination    as    already    described,   but 
omitting  the  addition  of  the  alkali  sulphate  (the  solution  contains  NaHS04 
from  the  thiosulphate). 

9.  It  is  evident  that  ammonium  salts  may  be  analyzed  for  ammonia  by 
simply  distilling  them  with  an  excess  of  alkali,  absorbing  the  ammonia  in 


VOLUMETRIC  ANALYSIS  123 

an  excess  of  standard  acid,  etc.  Moreover,  nitric  acid  and  nitrates  may 
be  quantitatively  reduced  to  ammonia  and  determined  in  this  way  (see 
Part  V,  Problems  62,,  64,  and  89). 

Certain  other  salts,  as  acetates,  may  be  analyzed  in  an  analogous  manner 
by  distillation  with  phosphoric  acid  in  excess,  the  distillate  being  collected 
in  standard  alkali  and  the  excess  of  the  latter  titrated,  with  the  use  of  phenol- 
phthalein. 


124  QUANTITATIVE  CHEMICAL  ANALYSIS 

B.    METHODS   OF  OXIDATION  AND  REDUCTION 

Standard  Solutions.  In  most  oxidation  and  reduction  pro- 
cesses the  standard  solution  actually  employed  in  the  titration 
is  one  of  an  oxidizing  reagent,  though  it  is  often  an  advantage 
also  to  have  available  a  suitable  standard  reducing  solution. 
Reducing  substances  are  nearly  always  determined  by  direct 
titration,  but  in  the  case  of  oxidizing  substances  it  is  often  more 
satisfactory  to  first  add  a  reducing  substance  in  excess,  and  then 
to  titrate  the  excess  of  this  reagent,  or  a  product  (as  iodine) 
liberated  from  it  by  the  oxidizing  substance. 

The  most  important  reagents  used  in  the  preparation  of  stand- 
ard oxidizing  solutions  are  potassium  permanganate,  iodine, 
potassium  dichromate,  potassium  bromate,  and  ferric  chloride; 
and  the  reagents  most  useful  for  standard  reducing  solutions  are 
ferrous  ammonium  sulphate,  oxalic  acid,  sodium  thiosulphate, 
sodium  arsenite,  and  titanous  chloride.  Other  substances 
are  frequently  used  in  volumetric  processes  for  purposes  of  oxi- 
dation and  reduction,  but  not  often  as  standard  solutions. 

In  the  titration  of  a  specific  reducing  substance,  such  as 
oxalic  acid,  it  is  of  course  necessary  to  employ  a  standard  solu- 
tion which  will  rapidly  and  completely  oxidize  the  substance; 
in  this  instance,  potassium  permanganate.  In  general,  the 
following  pairs  give  satisfactory  results :  potassium  permanganate 
and  ferrous  iron  or  oxalic  acid ;  iodine  and  sodium  thiosulphate 
or  arsenite;  potassium  dichromate  and  ferrous  iron;  ferric 
iron  and  titanous  chloride. 

Indicators.  With  respect  to  the  indicators  employed,  po- 
tassium permanganate,  owing  to  its  intense  coloring  power, 
is  its  own  indicator ;  the  slightest  excess  is  easily  visible  in  other- 
wise colorless  (or  even  in  certain  faintly  colored)  reaction  mix- 
tures. In  the  titration  of  ferrous  iron  with  dichromate  solu- 
tions, there  is  no  really  satisfactory  indicator  which  can  be 
added  to  the  solution ;  but  by  means  of  potassium  ferricyanide, 
as  an  outside  indicator,  it  is  possible  to  determine  with  great 


VOLUMETRIC  ANALYSIS  125 

accuracy  the  end-point  of  the  reaction.  In  the  case  of  iodine, 
starch  solution  is  employed  as  an  indicator.  The  use  of  these 
indicators  will  be  discussed  under  the  respective  processes. 

I.  DICHROMATE  PROCESSES 

Fundamental  Principles.  In  the  presence  of  acid,  ferrous- 
ion  is  quantitatively  oxidized  in  the  cold  to  ferric-ion  upon  the 
addition  of  potassium  dichromate  solution.  Since  hydrochloric 
is  the  most  efficient  acid  for  dissolving  the  ores  of  iron,  the  ti- 
tration is  almost  always  carried  out  in  the  presence  of  this 
acid: 

6  FeCl2+K2Cr207+i4  HC1=6  FeCl3+2  KCl-h2  CrCl3+7  H2O. 

^  The  end-point  of  the  titration  is  determined  by  means  of  a 
solution  of  potassium  ferricyanide.  As  this  point  is  approached, 
the  iron  solution  is  added  dropwise  to  the  indicator,  on  a  white 
test-plate  (or  sheet  of  paper),  and  the  mixture  examined  for  a 
blue  tint  (due  to  the  insoluble  compound  formed  by  the  inter- 
action of  ferrous-  and  ferricyanide-ion).  The  indicator  must  of 
course  be  sufficiently  dilute  to  be  almost  colorless,  and  must  not 
contain  any  ferrocyanide.  For  the  latter  reason,  since  the  solu- 
tion is  unstable,  it  must  be  prepared  just  before  use. 


THE  PREPARATION  AND  STANDARDIZATION  OF  THE  AP- 
PROXIMATELY TENTH-NORMAL  DICHROMATE  AND 
FERROUS  IRON  SOLUTIONS 

Procedure.  Dissolve  2.5  g.  of  potassium  dichromate  in  water, 
and  dilute  the  solution  to  500  cc. ;  also  dissolve  20  g.  of  ferrous 
ammonium  sulphate  and  5  g.  of  ammonium  sulphate  in  water, 
add  5  cc.  of  concentrated  sulphuric  acid,  and  dilute  to  500  cc. 
Thoroughly  mix  the  solutions,  allow  them  to  come  to  the  room 
temperature,  and  then  fill  a  burette  with  each,  with  care  to  ex- 
pel all  air  from  the  tips.    Dissolve  a  cubic-millimeter  fragment 


vO 


126  QUANTITATIVE  CHEMICAL  ANALYSIS 

of  pure  potassium  ferricyanide  in  15  cc.  of  water,  and  with  a 
glass  rod  transfer  drops  of  this  liquid  to  the  white  test-plate. 

Measure  out  into  an  Erlenmeyer  flask  about  20  cc.  of  the  iron 
solution,  add  10-15  cc.  of  6-normal  hydrochloric  acid  and  100 
cc.  of  water,  and  then  with  shaking,  run  inr  16-18  cc.  of  the  di- 
chromate  solution.  Now,  with  the  aid  of  a  glass  rod,  transfer  a 
drop  of  the  solution  to  a  drop  of  the  indicator,  being  careful  to 
wash  the  rod  before  returning  it  to  the  titration  mixture.  If  a 
blue  color  at  once  shows  up,  add  about  0.5  cc.  more  of  the  di- 
chromate,  and  test  again.  As  the  blue  tint  shows  up  more  and 
more  faintly,  add  smaller  and  smaller  amounts  of  the  dichromate 
solution,  and  use  larger  drops  of  the  titration  mixture  in  the 
tests,  until  finally  the  bluish  tinge  fails  to  materialize  in  30 
seconds.  Then,  for  'the  sake  of  comparison,  make  a  new  mix- 
ture beside  the  last  test ;  if  the  two  look  exactly  alike,  the  end- 
point  is  indicated.  While  it  is  best  not  to  overstep  the  end- 
point,  in  this  instance  this  can  do  no  serious  harm;  it  is  only 
necessary  to  add  more  of  the  ferrous  solution,  and  to  proceed 
with  the  titration. 

Repeat  this  operation  until  reliable  checks  are  obtained,  and 
from  the  (corrected)  data  calculate  the  value  of  the  ferrous  sul- 
phate in  terms  of  the  dichromate  solution ;  the  duplicates  should 
agree  within  two  parts  in  one  thousand. 

Standardize  the  dichromate  solution  as  follows:  Weigh  out 
three  0.14-0. 16  g.  portions  of  pure  iron  wire,  free  from  rust, 
bundhng  them  up  before  weighing  (handling  the  wire  with 
filter  paper)  so  that  they  may  easily  be  dropped  into  an  Erlen- 
meyer flask.  Treat  each  portion  as  follows :  Heat  20  cc.  of  6- 
normal  hydrochloric  acid  just  to  boihng  in  an  Erlenmeyer  flask, 
remove  the  flame,  and  drop  in  the  wire;  after  solution  is  ef- 
fected, boil  gently  for  2  minutes.  While  hot,  add  stannous 
chloride  solution,  drop  by  drop,  until  the  mixture  is  decolorized ; 
do  not  add  more  than  2  drops  in  excess.  Cool,  add  100  cc.  of 
cold  water,  and  then  pour  in  suddenly  with  shaking  15  cc.  of 
mercuric   chloride  solution.    After  an  interval  of   i   minute, 


VOLUMETRIC  ANALYSIS  127 

titrate  the  solution.  It  is  best  not  to  overstep  the  end-point; 
fewer  burette  readings  will  decrease  the  number  of  errors  in- 
volved. 

Notes.  —  i.  The  ionic  changes  due  to  oxidation  and  reduction  are  very 
marked  in  character,  and  the  electrical  charges  on  a  resultant  ion  may  even 
differ  in  sign  from  those  on  the  initial  ion ;  nevertheless,  in  the  equations 
which  represent  these  reactions,  the  algebraic  sum  of  the  charges  on  the 
one  side  is  always  equal  to  that  of  the  charges  on  the  other.  In  this  case, 
for  example,  we  have, 

6  Fe+++Cr207— +14  H+=6  Fe++++2  Cr++++7  H2O. 

2.  An  exactly  tenth-normal  solution  of  the  dichromate  may  be  prepared 
by  dissolving  2.4517  g.  of  the  pure  salt  in  water  and  accurately  diluting  the 

I  solution  to  500  cc.    The  commercial  salt  may  be  purified  by  recrystalliza- 
'    tion,  followed  by  drying  to  constant  weight  at  130°  (see  Part  I,  D). 

3.  The  presence  of  ammonium  sulphate  and  sulphuric  acid  in  the  standard 
iron  solution  seems  to  increase  the  stability  of  the  ferrous  condition. 

4.  The  iron  content  of  the  wire  purchased  for  this  standardization 
should  be  controlled  by  several  gravimetric  determinations.  The  wire  will 
stay  bright  if  kept  in  a  desiccator  over  solid  potassium  hydroxide,  but 
before  use  it  should  be  examined  for  rust ;  this  may  be  removed  by  means 
of  fine  emery  cloth. 

5.  The  short  boiling  after  solution  is  to  expel  any  hydrocarbon  that 
may  still  be  present  (the  iron  nearly  always  contains  some  carbon),  which 
otherwise  might  reduce  a  little  of  the  dichromate.  This  expulsion  is  even 
more  important  when  the  wire  is  used  as  a  standard  in  connection  with 
potassium  permanganate. 

6.  In  this  titration,  the  iron  to  be  determined  must  of  course  be  wholly 
in  the  ferrous  condition.  It  is  common  to  reduce  ferric  iron  to  ferrous  by 
means  of  stannous  chloride,  zinc  and  acid,  or  hydrogen  sulphide,  depending 
upon  the  conditions;  but  wherever  permissible,  it  is  most  convenient  to 
use  stannous  chloride.     (2  FeCl3+SnCl2=2  FeCl2+SnCl4.) 

Since  stannous  chloride  will  readily  reduce  dichromate  solution,  the 
excess  of  this  substance  must  be  oxidized  before  the  titration  of  the  iron. 
This  is  done  by  means  of  mercuric  chloride  (SnCl2+2  HgCl2=SnCl4 
-f  Hg2Cl2),  the  chlorides  of  mercury  being  without  influence  on  the  titration. 
If  present  in  any  quantity,  however,  stannous  chloride  is  likely  to  reduce 
the  mercury  partially  to  metal,  and  metallic  mercury  readily  reduces 
dichromate  solution.  (SnCl2-f-Hg2Cl2=SnCl4+2  Hg.)  Therefore,  the 
stannous  chloride  should  be  used  in  very  slight  excess,  and  to  this  end  it  is 
gradually  added  to  the  hot  ferric  solution,  before  the  dilution  of  the  latter. 


128  QUANTITATIVE  CHEMICAL  ANALYSIS 

On  the  other  hand,  the  excessive  reduction  of  the  mercury  is  less  likely  if  the 
stannous  chloride  is  cold  and  very  dilute,  and  if  the  mercuric  chloride  is 
suddenly  added  in  large  excess.  (The  formation  of  a  grayish  precipitate 
is  a  suflScient  indication  that  the  solution  should  be  discarded.) 

7.  The  dichromate  method  for  the  determination  of  iron  is  capable  of 
furnishing  very  accurate  values ;  but,  in  addition  to  the  usual  precautions, 
the  solution  should  be  titrated  promptly  after  the  precipitation  of  the  calomel 
(before  aeration  sets  in),  and  the  tests  for  the  end-point  should  not  be  begun 
until  that  point  is  fairly  imminent.  In  connection  with  the  determination 
of  this  end-point,  see  also  Note  2  under  the  gravimetric  determination  of 
iron. 


VOLUMETRIC  ANALYSIS  129 

THE  DETERMINATION  OF  IRON   IN  SIDERITE 

Procedure.  Weigh  out  0.23-0.25  g.  portions  of  the  powdered 
ore  into  Erlenmeyer  flasks,  moisten  the  samples  with  water, 
and  add  to  each  20  cc.  of  6-normal  hydrochloric  acid  and  about 
0.2  g.  of  potassium  chlorate.  Heat  not  quite  to  boiling  as  long 
as  there  is  solvent  action,  and  to  the  hot  solution  add  stannous 
chloride  solution,  drop  by  drop,  not  more  than  2  drops  in  ex- 
cess. Cool,  add  100  cc.  of  water,  shake,  and  then  pour  in  15  cc. 
of  mercuric  chloride  solution.  After  an  interval  of  i  minute, 
titrate  the  solution.  Report  the  percentage  of  iron  in  the 
ore. 

Notes.  —  i.  Siderite  (native  ferrous  carbonate)  is  an  iron  ore  plentiful 
in  England.  It  is  apt  to  contain  organic  matter,  and,  in  order  to  destroy 
this,  potassium  chlorate  is  added  during  its  solution. 

2.  Other  ores  can  of  course  be  analyzed  by  this  method.  Since  most  of 
them  contain  ferric  iron,  or,  since  in  the  case  of  ferrous  ores  the  iron  is  gen- 
erally oxidized  during  the  preparation  of  the  solution,  the  amount  of  stannous 
chloride  to  be  added  is  greater  than  that  required  in  the  standardization 
against  iron  wire.  In  no  case,  however,  should  the  excess  used  exceed  2 
drops. 

3.  For  a  more  general  method  of  dissolving  iron  ores,  see  the  determina- 
tion of  iron  by  means  of  potassium  permanganate,  and  also  Note  i  under 
that  method. 


I30  QUANTITATIVE  CHEMICAL  ANALYSIS 

II.  PERMANGANATE  PROCESSES 

Fundamental  Principles.  In  acid  solution,  potassium  per- 
manganate readily  oxidizes  ferrous-ion  to  ferric,  at  the  ordinary- 
temperature.  Also,  starting  at  80-90°,  it  reacts  quantitatively 
with  oxalic  acid,  which  it  oxidizes  to  carbonic  acid.  Though  in 
reality  the  reactions  are  not  so  simple,  the  quantitative  rela- 
tionships are  accurately  represented  by  the  following  equations : 

10  FeS04+2  KMn04+9  H2S04=5  Fe2(S04)3+2  KHSO4 

+2MnS04+8H20; 
and  5  H2C2O4+2  KMn04+4  H2S04=2  KHSO4+2  MnS04 

+10  CO2+8  H2O. 
Or,  more  simply  expressed, 

5  Fe++-hMn04-+8  H+=  5  Fe+++-|-Mn++  +4  H2O ; 
and        5  C2O4— +2  Mn04-+i6  H+=  2  Mn++-|-io  CO2+8  H2O. 

In  a  hot  neutral  or  faintly  acid  solution,  in  the  presence  of 
zinc  salts,  potassium  permanganate  oxidizes  manganous  salts 
quantitatively  in  the  sense  of  the  equation, 

3  Mn+++2  Mn04-+2  H20=4  H+-h5  MnOg. 

From  these  equations  it  is  readily  seen  that  for  use  in  acid 
solution  the  normal  weight  of  the  salt  is  one  fifth  of  a  mol,  or 
31.61  g.,  while  for  use  in  the  determination  of  manganese  the 
normal  weight  is  one  third  of  a  mol,  or  52.68  g.  It  is  not  cus- 
tomary, however,  in  volumetric  analysis,  to  employ  solutions  of 
permanganate  of  greater  than  tenth-normal  concentration. 

In  addition  to  the  above,  potassium  permanganate  is  capable 
of  oxidizing  stannous,  cuprous,  and  mercurous  salts,  antimonius, 
arsenious,  nitrous,  and  sulphurous  acids,  hydrogen  sulphide, 
ferrocyanides,  and  many  other  substances. 

Furthermore,  as  a  less  desirable  feature,  the  permanganate 
is  capable  under  certain  conditions  of  oxidizing  free  hydrochloric 
acid,  with  the  liberation  of  chlorine.  This  action,  though  al- 
most imperceptible  with  cold  dilute  hydrochloric  acid,  is  vigor- 


VOLUMETRIC  ANALYSIS  131 

ously  catalyzed  by  ferrous  iron.  With  suitable  modifications, 
however,  it  is  possible  to  obtain  very  exact  results  in  the  pres- 
ence of  hydrochloric  aid,  even  in  the  titration  of  iron ;  but,  other 
things  being  equal,  in  acid  solution,  it  is  preferable  to  carry  out 
permanganate  titrations  in  the  absence  of  chlorides. 

Even  at  tenth-normal  concentration,  potassium  permanganate 
solution  is  so  deeply  colored  that  the  lower  line  of  the  meniscus 
is  not  clearly  visible  in  an  ordinary  burette;  readings  must 
therefore  be  made  from  the  upper  edge.  This  disadvantage, 
however,  is  more  than  offset  by  the  fact  that  the  presence  of  a 
single  drop  in  excess,  in  an  otherwise  colorless  solution,  may  be 
recognized  with  great  ease ;  as  its  own  indicator,  it  leaves  nothing 
to  be  desired. 

The  permanganate  solution  should  not  be  placed  in  burettes 
with  rubber  connections ;  it  is  more  or  less  rapidly  reduced  by 
most  organic  substances. 

THE  PREPARATION  AND  STANDARDIZATION  OF  AN  AP- 
PROXIMATELY TENTH-NORMAL  SOLUTION  OF  POTAS- 
SIUM PERMANGANATE 

Procedure.  Dissolve  3.25  g.  of  the  salt  in  250  cc.  of  warm 
water,  allow  to  cool,  dilute  to  i  liter,  and  mix  thoroughly.  The 
value  of  this  solution  is  apt  to  change  slowly,  especially  at  first, 
and  for  this  reason  the  solution  should  be  allowed  to  stand  for 
several  days,  and  then  filtered  through  a  layer  of  asbestos  to  re- 
move the  precipitate  of  hydrated  manganese  dioxide.  After 
thorough  mixing,  it  is  then  ready  for  standardization.  The 
solution  should  be  preserved  in  glass-stoppered  bottles,  and 
should  be  protected  from  heat  and  light.  Thus  prepared  and 
preserved,  it  will  retain  its  oxidizing  value  for  months.  The 
solution  is  said  to  be  still  more  stable  if  it  is  made  very  slightly 
alkaline  with  potassium  hydroxide  (before  standardization,  of 
course). 

Weigh  out  accurately  into  700  cc.  Erlenmeyer  flasks  several 
0.1 2-0. 1 4  g.  samples  of  pure  sodium  oxalate,  previously  dried 


132  QUANTITATIVE  CHEMICAL  ANALYSIS 

at  IIO-I20®;  dissolve  each  sample  in  250  cc.  of  hot  water  (80- 
90°),  with  the  addition  of  30  cc.  of  6-normal  sulphuric  acid,  and 
titrate  at  once  with  the  permanganate  solution.  At  first,  the  per- 
manganate should  be  added  drop  by  drop,  with  shaking  after 
each  addition  until  the  color  disappears.  After  several  drops 
have  been  added,  the  solution  may  be  run  in  slowly  (10-12  cc. 
per  minute)  with  continuous  shaking.  Toward  the  end  of  the 
titration,  particular  care  must  be  taken  to  allow  the  color  due 
to  each  drop  to  disappear  before  the  addition  of  the  next,  in 
order  to  avoid  passing  the  end-point.  Titrate  to  the  first  per- 
manent pink.  The  temperature  at  the  end  of  the  titration  must 
not  be  below  60°. 

From  the  data  obtained,  calculate  the  normality  factor  of 
the  solution.  Duplicate  values  should  check  within  two  parts 
in  one  thousand. 

Notes.  —  i.  It  is  not  satisfactory  to  prepare  a  standard  solution  by 
directly  weighing  out  the  calculated  quantity  of  potassium  permanganate, 
even  after  the  latter  has  been  purified  by  recrystallization.  The  solution 
should  be  prepared  as  described  in  the  procedure,  and  then  standardized 
against  iron  wire  or  sodium  oxalate.  While  ferrous  ammonium  sulphate, 
oxalic  acid,  potassium  tetroxalate,  and  other  compounds  might  be  used  as 
standards,  iron  wire  and  sodium  oxalate  are  readily  obtainable  in  a  suffi- 
ciently pure  condition,  and  being  non-hygroscopic  and  free  from  water  of 
crystallization,  their  composition  is  less  subject  to  change. 

2.  Upon  treating  a  given  weight  of  pure  sodium  oxalate  with  an  excess 
of  sulphuric  acid,  the  corresponding  weight  of  oxalic  acid  is  set  free;  so 
that  the  use  of  this  salt  as  a  standard  merely  enables  us  easily  to  measure 
out  a  specific  amount  of  oxalic  acid.  The  oxidation  of  the  oxalic  acid  by 
the  permanganate  is  at  first  slow,  and  the  permanganate  should  be  added 
dropwise,  with  full  time  for  decolorization  between  successive  drops.  After 
a  certain  small  amount  of  manganous  sulphate  has  been  produced  in  the 
solution,  however,  the  speed  of  the  reaction  is  very  greatly  increased  (by 
the  catalytic  action  of  this  substance)  and  the  permanganate  may  be  run 
in  much  faster.  (See  R.  S.  McBride:  Jour,  Amer.  Chem.  Soc.y  vol.  34, 
p.  415  (1912).) 


VOLUMETRIC  ANALYSIS  133 

THE  DETERMINATION  OF  IRON  IN  HEMATITE 

Principles.  One  of  the  most  accurate  methods  for  the  deter- 
mination of  iron  is  based  upon  the  oxidation  of  a  chloride-free 
ferrous  sulphate  solution,  in  the  presence  of  sulphuric  acid,  with 
potassium  permanganate.  Under  these  conditions,  ferrous 
iron  is  oxidized  and  permanganate  is  reduced,  according  to  the 
equation : 

MnOr+S  Fe+++8  H+=Mn+++5  Fe++++4  H2O. 

But  if  chlorides  are  present,  some  of  the  permanganate  will 
be  reduced  by  these,  with  the  liberation  (and  partial  escape)  of 
chlorine,  and  the  results  will  be  somewhat  high : 

2  Mn04-+i6  H++10  Cl-=  2  Mn+++8  H2O+5  CI2. 

Upon  the  addition  of  the  permanganate  to  a  cold,  dilute  solu- 
tion of  hydrochloric  acid  alone,  or  to  one  containing  ferric  iron, 
no  chlorine  is  evolved;  ferrous  iron,  therefore,  seems  to  accel- 
erate this  reaction  by  catalysis. 

Nevertheless,  since  in  dissolving  iron  ores  it  is  nearly  always 
necessary  to  use  strong  hydrochloric  acid,  to  which  it  is  often 
well  to  add  a  Uttle  stannous  chloride,  and  since  stannous  chloride 
is  a  most  convenient  reagent  for  the  reduction  of  ferric  iron  to 
the  ferrous  condition,  it  is  desirable,  if  possible,  to  carry  out  the 
titration  in  the  presence  of  fairly  large  quantities  of  chlorides. 

Now  it  has  been  shown  that  if,  when  chlorides  are  present,  a 
small  quantity  of  manganous  salt  is  added  to  the  solution,  the 
ferrous  iron  alone  is  oxidized,  and  that  accurate  titrations  can 
be  performed  (Zimmermann).  But  the  end-point  is  somewhat 
indistinct,  owing  to  the  yellow  tint  of  the  ferric  chloride  pro- 
duced. This  difficulty  can  be  overcome  by  the  addition  of 
phosphoric  and  sulphuric  acids  (Reinhardt),  which  have  recently 
been  shown  to  combine  with  ferric  iron  to  form  colorless  com- 
plexes such  as  H[Fe(S04)2],  H3[Fe(P04)2],  and  H6[Fe(P04)3] 
(Weinland  and  Ensgraber,  Zeitschrift  fiir  anorganische  Chemie, 


134  QUANTITATIVE  CHEMICAL  ANALYSIS 

vol.  84,  p.  349) ;  the  large  excesses  of  these  acids  repress  the 
dissociation  of  these  complexes  and  insure  a  colorless  solution. 

Procedure.  Weigh  out  three  samples  of  the  finely  ground 
ore,  of  about  0.25  g.  each,  into  100  cc.  beakers.  To  each  sample 
add  15  cc.  of  6-normal  hydrochloric  acid  and  2  cc.  of  stannous 
chloride  solution,  and  gently  heat  the  covered  beakers  for  10-15 
minutes,  until  nothing  other  than  a  small,  white,  sandy  residue 
remains  undissolved.  If  the  hot  solution  is  at  all  yellow,  dis- 
charge this  color  by  adding  stannous  chloride  solution,  one  drop 
at  a  time,  with  stirring ;  avoid  an  excess  of  more  than  two  drops. 
If,  however,  after  the  heating,  the  solution  is  colorless,  stannous 
chloride  is  present  in  unknown  excess,  and  must  be  oxidized 
by  adding  permanganate  solution  (not  to  be  counted,  of  course, 
in  the  volume  required  for  the  titration)  drop  by  drop  with 
stirring,  until  the  yellow  color  due  to  ferric  iron  appears;  dis- 
charge this  color  as  above  directed,  with  stannous  chloride  so- 
lution, one  drop  in  excess. 

After  cooling,  dilute  the  colorless  solution  with  50  cc.  of  cold 
water,  and  transfer,  with  stirring,  to  a  700-cc.  beaker  containing 
10  cc.  of  mercuric  chloride  and  50  cc.  of  water.  (If,  instead  of  a 
white  precipitate  of  calomel,  a  gray  precipitate  of  mercury  is 
formed  at  this  point,  the  solution  must  be  discarded.)  Dilute 
the  mixture  with  cold  water  to  about  500  cc,  add  8-10  cc.  of 
the  Zimmermann-Reinhardt  solution,^  and  titrate  at  once  with 
the  standard  permanganate  solution.  Add  the  permanganate 
slowly,  with  constant  stirring,  finally  in  single  drops,  until  the 
pink  color  flashes  throughout  the  solution  and  persists  for  15-20 
seconds ;  do  not  pass  the  end-point.  Report  the  percentage  of 
iron  in  the  ore. 

Notes.  —  i.  Many  iron  ores  are  not  completely  decomposed  by  hydro- 
chloric acid,  the  insoluble  residue  containing  more  or  less  iron,  as  silicate, 
titaniferous  iron,  etc.     Unless  iron  is  known  to  be  absent  in  the  insoluble 


*  Made  by  dissolving  67  g.  of  MnS04  .  4  H2O  in  500  cc.  of  water,  adding  138  cc. 
of  phosphoric  acid  (sp.  gr.,  1.7)  and  130  cc.  of  sulphuric  acid  (sp.  gr.,  1.84),  and 
diluting  with  water  to  one  liter. 


VOLUMETRIC  ANALYSIS  135 

residue,  the  finely  ground  sample  should  be  digested  on  the  hot  plate  with 
10  cc.  of  hydrochloric  acid  until  the  residue  is  white,  or  until  there  appears 
to  be  no  further  action ;  if  the  ore  contains  carbonaceous  matter,  a  little 
potassium  chlorate  should  be  added.  Finally  evaporate  to  dryness,  ex- 
tract with  5  cc.  of  hydrochloric  acid,  dilute  with  10  cc.  of  water,  allow  to 
settle,  and  decant  the  clear  liquid  through  a  small  filter,  transferring  the 
residue  to  the  filter  and  washing  with  as  little  cold  water  as  possible.  Ignite 
the  filter  and  residue  in  a  small  platinum  crucible,  allow  to  cool,  and  add 
20-30  drops  of  sulphuric  acid  and  twice  as  much  hydrofluoric  acid.  Heat 
carefully,  and,  if  the  residue  is  dissolved,  evaporate  to  white  fumes,  allow 
to  cool,  dissolve  in  water,  and  add  to  the  solution  at  first  obtained.  If, 
however,  this  treatment  fails  to  decompose  the  residue,  drive  off  most  of 
the  sulphuric  acid,  add  0.5-0.6  g.  of  potassium  bisulphate,  and  heat  gradually 
until  the  bisulphate  is  quite  liquid  and  fumes  of  sulphuric  acid  are  given  off 
whenever  the  lid  of  the  crucible  is  raised.  When  the  black  specks  have 
disappeared,  allow  the  crucible  to  cool  and  dissolve  the  salt  in  the  crucible 
with  hot  water  and  a  few  drops  of  hydrochloric  acid. 

In  case  ferric  iron  has  been  dissolved  in  hydrochloric  acid  in  contact  with 
platinum,  the  solution  should  be  oxidized  with  bromine  water  and  the  iron 
precipitated  with  ammonia ;  i.e.  if  it  is  desired  to  use  stannous  chloride  in 
the  reduction.  The  ferric  hydroxide  can  then  be  redissolved  (after  washing 
it  with  hot  water)  in  hydrochloric  acid  and  reduced.  Otherwise  the  iron 
solution  will  contain  a  small  quantity  of  platinum,  4  FeCl3-l-2  HCl+Pt 
=  4  FeCl2+H2PtCl6,  which  gives  a  characteristic  ferric-iron  color  with 
stannous  chloride,  and  prevents  the  recognition  of  the  point  at  which  the 
iron  is  reduced. 

2.  Three  samples  should  be  taken,  in  order  that  one  may  be  used  for  a 
rapid  preliminary  titration.  Having  ascertained  in  a  rough  manner  the 
iron  content  of  the  sample,  the  final  titrations  are  greaty  facilitated. 

3.  Stannous  chloride  is  a  great  help  in  the  solution  of  many  ores  con- 
taining ferric  iron.  Apparently  the  diflStcultly  soluble  particles  of  hematite 
are  continuously  reduced  at  the  surface  to  ferrous  oxide,  which  is  much 
more  readily  dissolved  by  the  acid. 

4.  Some  common  agents  for  the  reduction  of  ferric  iron  are  zinc,  sulphur- 
ous acid,  and  hydrogen  sulphide ;  stannous  chloride  is  excluded  unless  the 
titration  is  to  be  made  by  the  Zimmermann-Reinhardt  method.  In  that 
case  it  should  be  carefully  added,  in  very  slight  excess,  to  the  hot,  concen- 
trated, acid  solution  (cf.  the  standardization  of  dichromate  solution.  Note  6). 

5.  Soluble  salts  of  mercurous  mercury  are  readily  oxidized  by  potassium 
permanganate  in  acid  solution.  Mercurous  chloride,  however,  is  exceed- 
ingly insoluble,  and,  provided  only  a  very  small  quantity  is  suspended  in 


136  QUANTITATIVE  CHEMICAL  ANALYSIS 

the  solution,  its  action  is  so  slow  that  the  end-point  of  the  titration  can  be 
accurately  fixed.  The  pink  color  which  flashes  throughout  the  solution 
at  the  end  of  the  titration  is,  however,  not  permanent,  and  for  that  reason 
the  time-Hmit  set  should  be  closely  observed.  For  the  greatest  accuracy, 
the  permanganate  should  of  course  be  standardized,  under  exactly  the  same 
conditions,  against  a  known  quantity  of  metallic  iron.  But  the  error  due 
to  the  use  of  a  solution  standardized  against  sodium  oxalate  is  for  most 
purposes  negligible. 

6.  For  a  rapid  method  for  the  reduction  of  ferric  iron  by  means  of  zinc, 
see  Notes  i  and  2  under  the  Determination  of  Phosphorus  in  Steel.  It 
should  be  noted  that  titanium  is  also  reduced  by  zinc,  but  not  by  the  other 
agents  mentioned ;  with  the  use  of  zinc,  therefore  the  presence  of  titanium 
would  lead  to  high  results. 


VOLUMETRIC  ANALYSIS  137 

THE  DETERMINATION  OF  CALCIUM  IN  LIMESTONE 

Procedure.  Instead  of  igniting  the  precipitate  of  calcium 
oxalate,  obtained  from  the  limestone  by  double  precipitation 
according  to  the  procedure  described  in  Part  II,  and  weighing 
it  as  calcium  oxide,  the  calcium  may  be  determined  volumetrically 
as  follows :  Wash  the  reprecipitated  calcium  oxalate  by  decanta- 
tion,  keeping  it  as  far  as  possible  in  the  precipitation  vessel,  and 
decompose  this  precipitate  by  slowly  pouring  through  the  filter 
at  least  six  5  cc.  portions  of  hot,  3-normal  sulphuric  acid,  wash- 
ing afterwards  with  hot  water,  and  receiving  the  acid  filtrate 
and  washings  in  the  beaker  containing  the  bulk  of  the  precipitate. 
Dilute  this  mixture  to  100  cc.  and  warm  gently,  with  stirring, 
to  completely  decompose  the  calcium  oxalate.  Allow  the  mixture 
to  cool,  transfer  it  quantitatively  to  a  250-cc.  measuring  fla^k, 
and  dilute  to  the  mark  with  water,  finally  mixing  the  solution 
by  pouring  it  into  a  clean  dry  beaker  and  back  into  the  flask. 

Measure  out  by  means  of  a  pipette  50.00  cc.  portions  of  this 
solution,  add  to  each  30  cc.  of  6-normal  sulphuric  acid,  dilute  to 
300  cc,  heat  to  90°,  and  titrate  as  already  described  with  the 
standard  permanganate  solution.  Remembering  that  only  one 
fifth  of  the  sample  was  used  in  each  titration,  calculate  the  per- 
centage of  CaO  in  the  limestone. 

Note.  —  The  reactions  involved  in  the  volumetric  determination  of 
calcium  are : 

CaC204+H2S04=CaS04+H2C204;  and  5  H2C204-h2  KMnO* 
+4  H2S04=  2  KHSO4-I-2  MnS044-io  COg+S  H2O. 

Therefore,  the  normal  weight  of  calcium  oxide  in  this  case  is  one  half  of  a 
mol ;  ox  0.1  N  permanganate  solution  has  a  calcium  oxide  value  of  0.00280  g. 
per  cubic  centimeter. 


138 


QUANTITATIVE  CHEMICAL  ANALYSIS 


THE  DETERMINATION  OF  THE  MNO2-VALUE  OF  PYROLUSITE 

Procedure.  Weigh  out  two  portions  of  the  finely  ground 
mineral,  of  about  0.3  g.  each,  into  700  cc.  Erlenmeyer  flasks. 
Calculate  the  weight  of  ferrous  ammonium  sulphate, 
Fe(NH4S04)2 .  6  H2O,  required  to  react  with  each  sample,  on 
the  basis  that  it  is  pure  manganese  dioxide: 

2  FeS04+Mn02+2  H2S04=Fe2(S04)3+MnS04+2  H2O, 

and  weigh  out  accurately  portions  of  the  pure  salt  0.15-0.20  g. 
in  excess  of  the  calculated  amounts,  into  the  corresponding 
flasks.  Add  to  each  flask  50  cc.  of  water  and  50  cc.  of  6-normal 
sulphuric  acid,  cover  the  flasks,  and  heat  to  boiling  until  the 
action  is  complete.  Finally,  dilute  to  about  300  cc,  and 
promptly  titrate  the  excess  of  ferrous  iron  with  the  standard 
permanganate  solution.  From  the  data  obtained,  calculate 
the  percentage  of  Mn02  in  the  sample. 

Notes.  —  i.  In  order  that  solution  shall  take  place  without  difficulty, 

the  mineral  should  be  ground  fine  enough  to  wholly  pass  through  a  loo-mesh 

sieve.    In  that  case,  a  moderate  excess  of  ferrous  iron  will  insure  rapid 

solution,  provided  the  mixture  is  not  diluted  before 

solvent  action  has  ceased. 

2.  A  solution  of  iron  wire  in  sulphuric  acid  may 
be  substituted  for  the  ferrous  ammonium  sulphate, 
but  in  that  case  there  is  more  danger  of  the  partial 
oxidation  of  the  iron  by  the  air.  For  example,  if  iron 
wire  is  used,  it  should  be  dissolved  in  sulphuric  acid 
out  of  contact  with  air,  and  the  air  should  not  have 
access  to  the  solution  during  cooling.  This  is  best 
accomplished  by  means  of  a  Contat-Gockel  valve, 
which  consists  of-  a  glass  bulb  with  an  inner  siphon,  as 
shown  in  the  figure.  In  the  bulb  is  placed  a  cold 
saturated  solution  of  sodiimi  bicarbonate,  through 
which  the  hydrogen  (and  steam)  evolved  in  the  flask 
bubbles.  After  all  the  iron  has  been  dissolved,  the 
liquid  is  boiled  for  a  few  minutes  longer,  and  the 
flame  is  removed.  As  the  flask  cools  off,  small  portions  of  the  bicarbonate 
are  at  intervals  sucked  into  the  flask  and  decomposed  by  the  acid  with  the 


VOLUMETRIC  ANALYSIS  139 

evolution  of  carbon  dioxide,  whereby  the  entrance  of  more  bicarbonate 
solution  is  prevented. 

For  other  methods  of  performing  this  analysis,  see  Part  V,  Problems 
23,  73,  and  74.  According  to  O.  L.  Barnebey  (/.  Ind.  Eng.  Chem.,  vol.  q, 
p.  961  (1917)),  the  use  of  oxalic  acid  in  place  of  the  ferrous  salt  yields  less 
reliable  results. 

3.  With  the  substitution  of  very  dilute  nitric  acid  for  sulphuric  acid 
in  the  above  procedure,  the  method  may  be  used  to  determine  the  PbOr 
value  of  red  lead,  or  minium,  Pb304,  and  of  lead  peroxide,  Pb02.  Of 
these  substances,  samples  of  i.o  and  0.8  g.,  respectively,  should  be  taken 
when  30-cc.  burettes  are  used.  It  is  better,  however,  to  make  use  of  an 
iodometric  method. 


140  QUANTITATIVE  CHEMICAL  ANALYSIS 

THE  DETERMINATION  OF  PHOSPHORUS  IN  STEEL 

Principle.  The  molybdic  anhydride  contained  in  ammonium 
phosphomolybdate,  (NH4)3P04 .  12  M0O3,  may  be  reduced  by 
zinc  in  the  presence  of  sulphuric  acid,  from  M0O3  to  M02O3; 
but  molybdenum  in  the  latter  condition  is  not  stable  in  the 
presence  of  air.  If,  however,  the  acidified  molybdate  solution 
is  passed  through  a  Jones  reductor  (see  below)  directly  into  a 
solution  of  ferric  sulphate,  the  sensitive  molybdate  compound  is 
oxidized  by  the  ferric  salt  with  the  formation  of  an  equivalent 
amount  of  ferrous  sulphate,  less  sensitive  to  the  atmospheric 
action.  The  molybdenum  solution  is  green  as  it  leaves  the 
reductor,  but  upon  mixing  with  the  ferric  salt  the  green  color 
disappears;  if  phosphoric  acid  is  added,  the  color  due  to  the 
presence  of  ferric  iron  is  destroyed.  The  decolorized  solution 
is  titrated  while  still  hot  with  tenth-normal  permanganate 
solution,  of  which  the  quantity  necessary  corresponds  to  the 
equation, 

5  M02O3+6  KMn04=3  K2O+6  MnO  +  io  M0O3. 

From  this  it  may  be  seen  that  5  P  =0=30  Mo203=o:36  KMn04<>9o  O, 
or  P=o=i8  O ;  one  cubic  centimeter  of  o.i  N  permanganate  solu- 
tion represents,  therefore,  0.0862  mg.  of  phosphorus. 

Procedure.  Weigh  out  two  samples  of  steel  drillings,  each 
sufficient  to  contain  1.7-2.0  mg.  of  phosphorus,  into  250-cc. 
Erlenmeyer  flasks.  Add  to  each  a  mixture  of  25  cc.  of  nitric 
acid  (sp.  gr.,  1.42)  and  75  cc.  of  water.  Suspend  in  the  neck 
of  each  flask  a  small  funnel  and  heat  until,  after  complete  solu- 
tion, the  oxides  of  nitrogen  have  been  expelled.  Dissolve  0.3- 
0.4  g.  of  KMn04  crystals  in  10  cc.  of  hot  water,  add  one  half  of 
this  solution  to  the  contents  of  each  flask,  and  boil  until  the 
permanganate  color  has  disappeared.  Remove  the  flame,  add 
sulphurous  acid  or  ammonium  bisulphite  solution,  a  few  drops 
in  excess,  to  dissolve  the  precipitated  oxides  of  manganese,  boil 
out  the  excess  of  sulphur  dioxide,  and  filter  the  solution;    re- 


VOLUMETRIC  ANALYSIS  141 

ceiving  the  filtrate  in  a  similar  flask.  Add  ammonia  to  the 
solution  with  stirring  until  a  permanent  precipitate  just  begins 
to  form,  and  then  add  nitric  acid  drop  by  drop  to  clear  up  the 
solution.  Finally,  at  a  temperature  of  40°,  add  40  cc.  of  molyb- 
date  solution,  close  the  flask  with  a  rubber  stopper,  and  shake 
vigorously  for  five  minutes;  allow  the  precipitate  to  settle. 
(At  this  point,  prepare  the  Jones  reductor  for  use,  as  described 
in  Note  2.) 

Now  filter  the  solution,  keeping  the  precipitate  as  far  as  pos- 
sible in  the  flask,  and  wash  by  decantation  with  a  solution  of 
ammonium  sulphate  acidified  with  sulphuric  acid  ^  until  the 
washings  give  no  test  for  molybdenum  with  ammonium  sulphide 
and  hydrochloric  acid.  Dissolve  the  precipitate  by  pouring 
through  the  filter  a  mixture  of  5  cc.  of  6-normal  ammonia  and 
20  cc.  of  water,  and  collecting  the  filtrate  and  washings  in  the 
precipitation  flask.  Acidify  the  solution,  which  should  have  a 
volume  of  about  60  cc,  with  10  cc.  of  sulphuric  acid  (sp.  gr., 
1.84)  and  promptly  pass  the  acidified  solution,  before  it  has  a 
chance  to  cool  off,  through  the  reductor  into  the  receiver  (collect- 
ing the  liquid  beneath  the  surface  of  100  cc.  of  a  solution  con- 
taining 25  g.  of  ferric  alum  and  40  cc.  of  sirupy  phosphoric  acid, 
sp.  gr.,  1.7,  per  liter),  preceded  by  100  cc.  of  hot  water  and  fol- 
lowed by  200  cc.  of  hot  dilute  sulphuric  acid  (i  140)  and  by  100  cc. 
of  hot  water.  See  that  no  air  enters  the  reductor  during  this 
entire  operation.  Titrate  the  reduced  solution  at  once  with 
tenth-normal  permanganate,  and  calculate  the  percentage  of 
phosphorus  in  the  steel  on  the  assumption  that  the  yellow  pre- 
cipitate contains  phosphorus  and  molybdenum  in  the  proportion 
indicated  by  the  formula  (NH4)3P04 .  12  M0O3. 

Notes.  —  i.  The  Jones  reductor,  which  also  is  useful  in  the  reduction 
of  ferric  iron  for  titration,  is  essentially  a  column  of  amalgamated  zinc, 
through  which  the  solution  is  passed  for  reduction.  It  is  assembled  as 
shown  in  the  figure  on  p.  142.    The  tube  A  is  of  about  18  mm.  inside 


1  Made  by  mixing  15  cc.  of  ammonia  (sp.  gr.,  0.90)  and  25  cc.  of  sulphuric  acid 
(sp.  gr.,  1.84)  with  one  liter  of  water. 


142 


QUANTITATIVE  CHEMICAL  ANALYSIS 


fSOmah 


diameter,  and  (for  this  reduction)  400  mm.  long ;  the  small  extension  tube 
is  of  6  mm.  inside  diameter  and  extends  downwards  300  mm.  from  the  stop- 
cock. The  outlet  of  the  tube  A  is  covered  with  a  layer  of  glass  beads; 
these  are  surmounted  by  a  small  plug  of  glass  wool ;  and  upon  this  is  placed 
a  layer  of  asbestos  felt,  not  exceeding  i  mm.  in  thickness.  Amalgamated 
zinc  is  then  added  to  within  50  mm.  of  the  top  of  the  tube,  and  this  may  be 

covered  with  glass  wool  and  as- 
bestos felt,  to  retain  any  soHd 
matter.  The  reductor  is  con- 
nected as  shown  with  the  suction 
flask  F,  which  in  turn  is  connected 
with  the  safety  bottle  Z),  to  avoid 
the  intake  of  any  contamination 
from  the  suction  apparatus. 

The  amalgamated  zinc  is  pre- 
pared by  dissolving  5-6  g.  of 
mercuric  chloride  in  250  cc.  of 
water,  in  a  i  liter  bottle,  with  the 
addition  of  5-10  cc.  of  6-normal 
hydrochloric  acid,  adding  to  this 
solution  500  g.  of  (18  mesh)  granu- 
lated zinc,  and  shaking  vigorously 
for  at  least  a  minute ;  the  liquid 
is  then  decanted,  and  the  zinc  thor- 
oughly washed  with  water. 

2.  For  use,  the  reductor  is  connected  as  shown  with  the  suction  appara- 
tus ;  and,  with  the  stopcock  almost  closed,  it  is  filled  with  warm  sulphuric 
acid  (25  cc.  of  the  concentrated  acid  in  i  liter) ;  the  stopcock  is  then  opened 
to  allow  the  acid  to  run  through  slowly.  Acid  is  poured  in  until  200-300  cc. 
have  passed  through,  and  then  with  liquid  still  in  the  funnel,  the  cock  is 
closed.  (In  using  this  apparatus,  see  that  no  air  enters  the  reductor; 
otherwise  hydrogen  peroxide  may  be  formed  from  oxygen  and  nascent 
hydrogen  and  vitiate  the  results.)  The  first  filtrate  is  rejected,  200  cc.  more 
of  the  warm  acid  are  passed  through,  followed  by  100  cc.  of  warm  water, 
and  to  this  liquid  (300  cc.)  the  standard  permanganate  solution  is  added,  one 
drop  at  a  time,  from  a  burette,  in  order  to  determine  the  volume  required 
to  color  the  acid  solution  alone.  This  amount  (required  by  impurities  from 
the  zinc)  must  be  subtracted  from  the  volume  required  in  the  subsequent 
titration. 

3.  Upon  dissolving  the  steel  in  nitric  acid  of  the  strength  indicated,  the 
phosphorus  is  gxidized,  and  none  of  it  is  lost  by  evolution  as  phosphine. 


VOLUMETRIC  ANALYSIS  143 

The  permanganate  is  subsequently  added  in  order  to  insure  the  complete 
oxidation  of  carbonaceous  matter  and  of  the  phosphorus  to  phosphoric  acid. 

4.  The  higher  oxides  of  manganese,  as  Mn02,  are  not  soluble  in  nitric 
acid.  Upon  the  addition  of  a  reducing  agent,  however,  such  as  hydrogen 
peroxide  or  sulphurous  acid,  their  solution  is  effected : 

MnOa+HaSOs  =  MnO+H2S04=  MnSOi+HsO. 

5.  In  connection  with  the  precipitation  of  phosphoric  acid  as  ammonium 
phosphomolybdate,  the  student  should  consult  the  notes  under  the  Deter- 
mination of  Phosphoric  Anhydride. 

6.  Since  the  molybdenum  in  the  precipitate  prepared  from  one  gram  of 
a  steel  containing  0.15%  of  phosphorus  would  require  by  this  method 
17.44  cc.  of  o.i-normal  permanganate  solution,  it  is  readily  seen  that  the 
process  is  a  rapid  one  for  arriving  at  very  accurate  results.  This  is  es- 
pecially true  if  the  permanganate  has  been  standardized  under  the  same 
conditions  against  a  steel  of  accurately  known  phosphorus  content;  in 
such  a  case,  it  would  be  unnecessary  to  correct  for  the  small  amount  of  iron 
extracted  from  the  (impure)  amalgamated  zinc,  since  this  would  be  the 
same  in  both  standardization  and  analysis. 

7.  It  scarcely  needs  to  be  pointed  out  that  the  method  is  not  suitable  for 
determining  phosphorus  or  phosphoric  acid  in  substances  containing  them 
in  large  amount.  This  would  require  for  titration  relatively  enormous 
quantities  of  permanganate  solution,  and,  what  is  still  worse,  it  would  be 
practically  impossible  to  completely  reduce  the  molybdenum.  K  small 
aUquot  portions  were  taken  for  reduction  and  titration,  any  error  of  meas- 
urement would  be  multiplied  by  a  very  large  factor  in  the  calculation  of 
the  result. 

8.  The  following  method  is  suitable  for  the  volumetric  determination  of 
phosphorus  or  phosphoric  acid  when  these  are  present  in  larger  amounts. 
The  phosphorus  or  phosphoric  acid  is  converted  into  ammonium  phospho- 
molybdate ;  this,  after  washing  with  KNO3  solution,  is  dissolved  in  an  excess 
of  standard  sodium  hydroxide  solution ;  and  the  resulting  solution  is  titrated 
with  standard  nitric  acid,  with  phenolphthalein  as  an  indicator.  Needless 
to  say,  the  sodium  hydroxide  should  be  standardized  under  identical  condi- 
tions against  a  sample  of  accurately  known  phosphorus  content. 


144  QUANTITATIVE  CHEMICAL  ANALYSIS 

THE  DETERMINATION  OF  MANGANESE   IN  AN  ORE 

Fundamental  Principles.  When  potassium  permanganate  is 
added  to  a  hot,  neutral,  or  very  faintly  acid  solution  of  manganese 
sulphate  an  action  takes  place  in  which  the  manganous  oxide 
of  the  sulphate  is  oxidized  at  the  expense  of  the  anhydride  of  the 
permanganate,  with  the  precipitation  of  hydrated  intermediate 
oxides  in  varying  proportions.  These  are  manganous  acid, 
MnO(OH)2,  and  hydrated  salts  of  manganous  acid.  The  es- 
sential changes  in  the  state  of  oxidation  may  be  represented  as 
follows : 

Mn207+3  MnO=5  MnOa; 

MnaOy+S  MnO=  5(Mn02 .  MnO) ; 
Mn207+i3  MnO=5(Mn02 .  2  MnO). 

It  is  clear,  then,  that  in  this  form  the  action  cannot  furnish  the 
basis  for  a  satisfactory  volumetric  method. 

It  has  been  found,  however,  that  under  suitable  conditions, 
in  the  presence  of  zinc-ion,  a  hydrated  manganite  of  zinc  is  pre- 
cipitated, which,  while  variable  in  composition,  contains  all  the 
manganese  in  the  quadrivalent  condition.  Thus  regulated,  the 
reaction  furnishes  a  valuable  means  for  the  determination  of 
manganese  (Volhard's  Method).  Although  the  composition  of 
the  precipitate  varies,  the  course  of  the  reaction  is  typically 
represented  by  the  following  equation : 

4  KMn04+5  ZnS04+6  MnS04+i4  H20=  2  K2SO4+9  H2SO4 

.      ,  -}-5  ZnfoMn/^^^  , 

or,  more  simply,  ^      \         n^O    J 2. 

2  Mn04-+3  Mn++-F2  H20=4  H+-f5  Mn02. 

Procedure.  Weigh  out  into  a  500-cc.  Erlenmeyer  flask  a 
sufiicient  quantity  of  the  very  finely  ground  ore  to  contain  about 
0.20  g.  of  manganese ;  add  3  g.  of  potassium  chlorate  and  20  cc. 
of  1 2 -normal  hydrochloric  acid,  and  boil  until  the  ore  is  com- 
pletely decomposed  and  the  chlorine  expelled.  Dilute  with 
water  to  about  50  cc. ;  quantitatively  transfer  the  cold  solution 


VOLUMETRIC  ANALYSIS  I45 

to  a  loo-cc.  measuring  flask;  dilute  to  the  mark  with  water; 
and  mix  thoroughly  by  pouring  the  contents  of  the  flask  into 
a  clean,  dry  beaker,  and  back  into  the  flask. 

Now,  from  a  burette  or  pipette,  measure  into  500-cc.  Erlen- 
meyer  flasks  four  20.00  cc.  portions  of  this  solution,  and  treat 
each  as  follows :  Dilute  with  water  to  100  cc,  heat,  and  to  the 
acid  solution  add  with  shaking  an  aqueous  suspension  of  zinc 
oxide,^  in  small  portions,  until  the  iron  is  completely  precipitated 
as  ferric  hydroxide ;  this  point  may  be  recognized  by  the  sudden 
coagulation  of  the  precipitate,  upon  shaking,  and  the  decoloriza- 
tion  of  the  brownish  colored  solution.  The  precipitate  should 
not  be  light  yellow,  but  should  have  the  characteristic  brownish 
red  color  of  ferric  hydroxide,  and  the  least  possible  excess  of 
zinc  oxide  should  be  used.  (Should  the  ore  contain  a  quantity 
of  iron  insufficiently  in  excess  of  that  required  by  any  phosphoric 
and  arsenic  acids  present,  then  5  cc.  of  a  solution  containing 
20  g.  of  ferric  chloride  per  Hter  should  be  added  before  the  pre- 
cipitation with  zinc  oxide.)  If  too  much  zinc  oxide  is  added, 
the  solution  will  be  milky ;  in  that  case  very  dilute  hydrochloric 
acid  should  be  added  drop  by  drop  to  the  hot  solution  until  the 
supernatant  liquid  just  becomes  clear. 

Finally  dilute  the  solutions  to  300  cc.  and,  at  80°,  treat  them 
successively  as  follows :  Run  into  the  first  solution  the  standard 
permanganate  in  5  cc.  portions,  until  after  continued  shaking 
the  liquid  retains  a  permanent  pink  tinge,  —  say  after  the  addi- 
tion of  the  fifth  portion  {i.e.  25  cc.) ;  into  the  second  solution 
run  5  cc.  less  permanganate  than  the  volume  previously  used 

1  Dissolve  100  g.  of  crystallized  zinc  sulphate  in  300  cc.  of  hot  water,  and  with 
stirring  cautiously  add  to  the  clear  solution  a  few  drops  of  a  solution  made  by 
dissolving  25-27  g.  of  pure  sodium  hydroxide  in  150  cc.  of  water,  until  the  zinc 
solution  remains  distinctly  turbid ;  then  add  a  little  bromine  water,  heat,  and  filter. 
To  the  filtrate  add  the  bulk  of  the  sodium  hydroxide  solution,  and  stir.  Rinse  the 
mixture  into  a  one-liter  bottle,  and  fill  the  latter  with  water.  The  mixture  should 
be  well  shaken  when  used.  (This  suspension  should  not  react  alkaline  with  phenol- 
phthalein,  and  a  10  cc.  portion  of  the  mixture,  when  cleared  up  with  sulphuric  acid, 
diluted  to  100  cc,  and  treated  with  one  drop  oi  0.1  N  KMn04,  should  be  perma- 
nently colored  pink.) 


146  QUANTITATIVE  CHEMICAL  ANALYSIS 

{e.g.  20  cc),  shake  until  the  pink  color  disappears,  and  then  finish 
the  titration  by  the  further  addition  of  permanganate  in  portions 
of  I  cc.  until  the  pink  color  persists  after  protracted  shaking, 
say  after  23.0  cc.  in  all  have  been  added;  to  the  third  solution 
add  at  once  i  .0  cc.  less  permanganate  than  the  total  volume  used 
in  the  second  case  (e.g.  22.0  cc),  and  continue  the  titration  with 
the  addition  of  0.20  cc.  portions,  until  the  hot  solution  matches 
in  color  a  solution  prepared  by  the  addition  of  o.io  cc.  of  the 
permanganate  to  300  cc.  of  water.  With  the  fourth  solution, 
repeat  this  titration.  If,  for  example,  22.60  cc.  of  the  perman- 
ganate have  been  used  in  each  of  the  last  two  titrations,  then 
this  quantity  minus  the  o.io  cc.  of  the  solution  used  for  com- 
parison should  be  taken  as  the  volume  actually  required.  Report 
the  percentage  of  manganese  in  the  ore. 

Notes.  —  i.  In  case  the  treatment  with  hydrochloric  acid  and  potas- 
sium chlorate  should  be  insufficient  to  thoroughly  decompose  the  ore  (in- 
dicated by  the  presence  of  a  dark-colored  residue),  the  residue  should  be 
filtered  off,  washed,  dried,  and  ignited  in  a  platinum  crucible.  It  should 
then  be  fused  with  sodium  carbonate,  the  melt  dissolved  in  hydrochloric 
acid,  and  the  solution  evaporated  to  dryness  in  a  porcelain  dish,  to  dehy- 
drate the  silica.  The  final  residue  should  be  moistened  with  hydrochloric 
acid,  taken  up  in  water,  and  filtered  into  the  Erlenmeyer  flask  contain- 
ing the  acid  filtrate  from  the  original  residue.  The  resulting  solution, 
which  contains  all  the  manganese,  is  then  evaporated  to  a  small  volume, 
transferred  to  the  measuring  flask,  and  treated  as  described  in  the 
procedure. 

2.  Upon  the  addition  of  zinc  oxide  to  the  acid  solution  of  the  ore,  the 
zinc  oxide  first  neutralizes  the  acid  with  the  formation  of  zinc  chloride,  and 
then  precipitates  the  iron  with  the  further  formation  of  zinc  chloride,  accord- 
ing to  the  reaction : 

2  FeCl3+3  ZnO-l-3  H20=  2  Fe(OH)3-|-3  ZnClz. 

In  this  way  sufficient  zinc-ion  is  introduced  into  the  solution  to  insure 
the  conversion  of  the  manganese  into  the  hydrated  manganite  of  zinc. 

3.  Although  it  is  often  recommended  to  convert  the  chlorides  in  the 
solution  into  sulphates  before  the  addition  of  zinc  oxide,  this  treatment  is 
not  necessary.  The  titration  of  manganese  in  a  dilute  neutral  solution 
with  potassium  permanganate  is  a  very  different  thing  from  that  of  ferrous 


VOLUMETRIC  ANALYSIS  147 

iron  in  a  dilute  acid  solution  containing  chlorides.  In  the  latter  case,  the 
ferrous  iron  catalyzes  the  reaction, 

2  KMn04+i6  HC1=  2  KCH-2  MnCU+S  CI2+8  HA 

some  of  the  chlorine  escapes,  and  there  is  a  tendency  to  high  results.  In 
the  former  case,  however,  nothing  is  present  in  the  solution  to  catalyze 
the  reaction  between  the  permanganate  and  the  small  quantity  of  hydro- 
chloric acid  which  is  formed ;  and,  although  the  solution  is  hot,  its  acid 
concentration  is  so  low  that  there  is  no  danger  from  this  source.  Start- 
ing with  0.2000  g.  of  an  ore  containing  20%  of  manganese,  for  example, 
the  total  quantity  of  acid  formed  in  the  titration  (e.g.  4  KMn044-S  ZnCl2 
+6  MnCla-f  14  H20=4  KCI+18  HCl+s  ZnO  .  Mn203(OH)2)  weighs  about 

L — X0.04,  or  somewhat  less  than  o.i  g. ;  and  this  quantity  in  a  volume 
Mn 

of  over  300  cc.  would  give  an  add  strength  of  less  than  o.oi-normal. 

4.  The  titration  should  be  performed  at  80-85°,  and  especial  care  should 
be  taken  not  to  heat  the  solution  too  hot. 

5.  For  the  greatest  accuracy,  in  spite  of  all  that  has  been  said  above, 
the  permanganate  solution  should  be  standardized  against  a  known  quantity 
of  manganese,  weighed  as  MnS04,  under  conditions  similar  to  those  to  be 
used  in  the  analysis. 


148  QUANTITATIVE  CHEMICAL  ANALYSIS 

III.  lODOMETRIC  PROCESSES 

Fundamental  Considerations.  Analyses  which  are  based  upon 
the  volumetric  determination  of  specific  quantities  of  iodine 
are  classed  under  the  head  of  iodometric  processes.  In  these 
processes,  either  a  standard  solution  of  iodine  is  used  to  effect 
a  definite  chemical  action,  or  the  iodine  liberated  in  a  reac- 
tion is  volumetrically  determined  by  titration  with  a  suitable 
solution. 

The  estimation  of  iodine  with  standard  thiosulphate  solution, 
with  the  use  of  starch,  is  a  very  exact  volumetric  process ;  the 
titration  may  be  performed  in  neutral  or  sHghtly  acid  solution. 
Moreover,  if  the  iodine  has  previously  been  liberated  from  an 
excess  of  hydriodic  acid  by  means  of  a  specific  oxidizing  agent, 
then  the  results  of  the  titration  may  be  expressed  directly  in  terms 
of  this  oxidizing  agent  itself.  For  example,  the  chromium  in  a 
sample  of  ore  may  be  converted  into  the  hexivalent  condition, 
and  subsequently  determined  on  the  basis  of  the  reactions : 

K2Cr207+6  KI-I-14  HC1=8  KCI+2  CrCla+y  H2O+3  I2; 

and  I2+2  Na2S203=  2  NaI+Na2S406. 

Or,  more  simply  expressed, 

Cr207— +6 1-+14  H+=  2  Cr++++7  H2O-I-3  I2 ; 

and  I2  +2S2O3— =2l--|-S406~. 

As  an  oxidizing  agent,  iodine  may  act  directly,  as  in  the  re- 
action with  sodium  thiosulphate,  or  it  may  act  indirectly  through 
hypoiodous  acid,  which  it  gives  with  water,  as  in  the  equations : 

l24-HOH±;:HI+HOI; 
and  Na2HAs03+HOI±;HH-Na2HAs04. 

It  will  be  seen  from  these  equations  that  a  one-tenth  normal 
solution  of  iodine  contains  one  tenth  of  a  gram-atom,  or  12.692 
g.  of  available  iodine  per  liter. 
It  is  of  interest  to  mention  that  while  chlorine  and  bromine 


VOLUMETRIC  ANALYSIS  149 

oxidize  sodium  thiosulphate  partially  to  sulphate,  iodine,  under 
analytical  conditions,  oxidizes  it  wholly  to  tetrathionate. 

The  solubility  of  iodine  in  water  is  too  small  for  the  prepara- 
tion of  even  a  one  tenth  normal  solution.  In  the. presence  of 
sufficient  potassium  iodide,  however,  the  iodine  dissolves  much 
more  readily,  owing  to  the  formation  of  an  unstable  but  soluble 
polyiodide  of  the  formula  KI3 : 

KH-I2 1^  KI .  I2,  or  I--}-l2 1^  (I .  I2)- 

In  the  presence  of  reducing  agents  iodine  is  removed  from  this 
equilibrium  mixture,  the  reaction  runs  to  completion  from  right 
to  left,  and  the  solution  can  be  used  as  though  it  were  a  simple 
solution  of  iodine.  The  potassium  iodide  used  in  the  prepara- 
tion of  the  solution  should  weigh  about  1.5  times  as  much  as 
the  iodine.  Moreover,  the  presence  of  potassium  iodide  in  the 
solution  renders  it  possible  to  employ  commercial  iodine  (which 
is  apt  to  contain  chlorine  as  an  impurity)  in  the  preparation  of 
the  standard  solutions;  the  chlorine  is  removed  according  to 
the  equation,  ICl-hKI=KCH-l2. 

In  performing  iodometric  titrations  in  the  presence  of  sul- 
phuric acid,  particular  attention  should  be  given  to  the  main- 
tenance of  suitable  analytical  conditions.  If,  for  example,  it  is 
desired  to  determine  copper  by  titrating  the  iodine  liberated  in 
the  reaction, 

2  CUSO4+2  H2SO4+4  KI:^Cu2l2+4  KHSO4+I2, 
it  is  not  sufficient  to  simply  add  potassium  iodide  and  sul- 
phuric acid  in  (unknown)  excess.  It  must  -  be  remembered 
that  such  a  mixture  will  contain  both  sulphuric  and  hydriodic 
acid,  and  that  if  the  concentration  of  either  is  too  great,  or  if 
the  solution  is  allowed  to  become  at  all  warm,  the  determination 
is  very  apt  to  be  spoiled : 

H2S04-h2  HI=H2S03-i-H20-hl2; 
H2S04+6HI=S+4H20-i-3l2; 
or,  in  extreme  cases, 

H2SO4+8  HI=H2S-f4  H2O-I-4  I2. 


I50  QUANTITATIVE  CHEMICAL  ANALYSIS 

Other  things  being  equal,  when  acid  solutions  are  required,  it  is 
better  to  use  acetic  acid  or  dilute  hydrochloric  acid. 

In  direct  titrations  with  iodine,  e.g.  in  the  presence  of  sodium 
bicarbonate,  it  is  best  to  work  in  the  absence  of  ammonium 
salts.  Such  solutions  are  very  faintly  alkaline,  especially  if  at 
all  warm;  and  in  the  presence  of  ammonium  salts  ammonia 
is  apt  to  be  liberated,  which  is  not  entirely  without  influence 
upon  the  titration. 

Iodine  solutions  should  not  be  placed  in  burettes  with  rubber 
stopcocks. 

Determination  of  the  End-point.  A  single  drop  of  one-tenth 
normal  iodine  solution  imparts  a  distinct  tint  to  200  cc.  of  water, 
and  in  many  titrations  with  this  solution  no  other  indicator  is 
required.  If,  however,  the  solution  to  be  titrated  contains  colored 
substances,  or  if  the  greatest  possible  accuracy  is  demanded,  a 
solution  of  starch  should  be  used  as  an  indicator.  Under  the 
proper  conditions,  the  presence  of  one  part  of  free  iodine  in 
several  milKons  of  solution  can  be  recognized  with  this  indi- 
cator, but  the  sensitiveness  of  the  reaction  and  the  color  pro- 
duced are  affected  by  a  number  of  factors.  The  test  is  decidedly 
more  sensitive  when  the  concentration  of  iodide-ion  (and  of 
hydrogen-ion  ^)  is  not  too  low,  and  when  the  quantity  of  starch 
present  is  sufficient  to  give  a  deep  blue  color. 

Under  less  favorable  conditions,  the  starch  may  give  a  greenish 
or  a  reddish  color ;  or  it  may  be  very  unreliable,  as  in  solutions 
containing  an  abnormally  low  iodide-ion  concentration.  How- 
ever, since  the  standard  iodine  solution  always  contains  potas- 
sium iodide,  and  since  an  iodide  is  always  one  product  of  the 
titration,  there  is  ordinarily  not  much  danger  from  this  source. 
Attention  should  be  directed  mainly  toward  the  observance  of 
uniform  conditions  in  all  related  titrations :  the  volumes  of  the 

*  Like  iodine,  the  blue  iodo-starch  is  incapable  of  existence  in  the  presence  of 
hydroxide-ion  in  appreciable  quantity,  and  titrations  therefore  should  not  be 
attempted  in  the  presence  of  alkali  hydroxides  or  carbonates.  In  cold,  dilute 
solution,  however,  alkali  bicarbonates  have  no  bad  influence. 


VOLUMETRIC  ANALYSIS  151 

solutions  titrated  should  be  approximately  equal,  the  starch 
solution  should  be  properly  prepared,  and  the  same  quantity  of 
it  should  be  added  for  each  titration.  Finally,  all  titrations 
should  be  made  in  the  cold ;  the  iodo-starch  blue  is  discharged 
by  heat. 

Preparation  of  the  Starch  Solution.  Rub  i  g.  of  powdered 
starch  with  a  Uttle  cold  water  to  a  thin  paste,  and  slowly 
add  this  to  200  cc.  of  boiling  water,  stirring  the  hot  mixture 
until  an  almost  clear  solution  is  obtained.  Allow  this  to  settle, 
and  filter  off  the  liquid.  Use  5  cc.  of  this  liquid  for  each 
titration. 

A  "  soluble  starch  ''  which  is  in  the  market  is  more  convenient, 
since  with  it  filtration  is  unnecessary.  A  solution  made  by  add- 
ing 200  cc.  of  boiling  water  to  i  g.  of  this  starch,  previously 
mixed  with  a  little  cold  water,  serves  the  purpose  well.  Use 
5  cc.  of  this  solution  for  each  titration. 

In  either  case,  the  starch  solution  should  be  freshly  prepared. 
If  a  great  many  titrations  are  to  be  made,  however,  it  is  ad- 
visable to  prepare  a  Hter  of  the  starch  solution;  a  number  of 
small  (50-100  cc.)  bottles,  filled  with  this  solution,  should  be 
heated  for  at  least  an  hour  in  boiling  water,  and,  while  still  in 
the  bath,  they  should  be  sealed  with  paraffined  stoppers  of  soft 
cork.  Thus  sterilized  and  preserved,  the  solution  will  remain 
sensitive  for  many  months.  After  a  bottle  has  been  opened, 
mold  is  likely  to  form  within  a  few  days. 

THE  PREPARATION  AND  STANDARDIZATION  OF  APPROX- 
IMATELY ONE  TENTH  NORMAL  SOLUTIONS  OF  IODINE 
AND  SODIUM  THIOSULPHATE 

Procedure.  Weigh  out  on  the  rough  balance  6.3-6.4  g.  of 
commercial  iodine,  add  it  to  a  solution  of  9  g.  of  potassium 
iodide  in  25  cc.  of  water,  in  an  Erlenmeyer  flask,  and  agitate 
the  mixture  until  the  iodine  is  completely  dissolved.  Dilute 
the  solution  to  500  cc,  in  a  measuring  flask,  and  mix  it  thor- 
oughly. 


/ 

152  QUANTITATIVE  CHEMICAL  ANALYSIS 

Heat  600-700  cc.  of  distilled  water  in  a  large  flask  and 
boil  for  about  5  minutes.  Stopper  the  flask  loosely,  and  allow 
the  water  to  cool.  Weigh  out  12.5  g.  of  sodium  thiosulphate, 
Na2S203  .  5  H2O,  introduce  it  into  a  500-cc.  measuring  flask, 
and  dissolve  it  in  about  200  cc.  of  the  cold,  freshly  boiled  water. 
Finally  dilute  to  the  mark  with  more  of  the  same  water,  a«d 
mix  thoroughly. 

After  these  solutions  have  come  to  the  room  temperature, 
fill  a  burette  with  each  (see  p.  150),  and  measure  out  20  cc.  of 
the  thiosulphate  solution,  into  an  Erlenmeyer  flask.  Add  to  this 
125  cc.  of  water  and  5  cc.  of  starch  solution,  and  then,  with 
shaking,  run  in  the  iodine  solution  until  the  faintest  blue  tinge 
persists  in  the  mixture.  If  the  end-point  is  overstepped,  ti- 
trate back  with  the  thiosulphate  solution.  {All  waste  solutions 
containing  iodine  and  potassium  iodide  should  he  poured  into  the 
vessel  provided  for  iodine  residues.)  From  the  (corrected)  data, 
calculate  the  ratio  of  the  two  solutions,  repeating  the  titration 
until  satisfactory  checks  are  obtained. 

Weigh  out  into  500  cc.  Erlenmeyer  flasks  two  o.i 2-0.13  g- 
portions  of  arsenious  oxide,  and  in  each  case  warm  gently  with 
10  cc.  of  6-normal  sodium  hydroxide.  To  the  resulting  solution 
add  50  cc.  of  water  and  2  drops  of  methyl  orange,  and  then 
cautiously  add  6-normal  hydrochloric  acid,  2  or  3  drops  in  ex- 
cess. Finally  add  75  cc.  of  a  cold  solution  of  sodium  bicar- 
bonate (made  by  dissolving  10  g.  of  the  pure  salt  in  150  cc. 
of  water)  and  5  cc.  of  starch  solution,  and  then  titrate  with 
the  iodine.  Do  not  overstep  the  end-point.  From  the  data 
obtained,  calculate  the  normality  factor  of  the  iodine. 
(Duplicate  values  should  agree  within  two  parts  in  one  thou- 
sand.) Also  calculate  the  normahty  factor  of  the  thiosulphate 
solution. 

Notes.  —  i.  Iodine  solutions,  like  chlorine  water,  are  acted  upon  by 
sunlight ;  and,  also,  iodine  is  readily  volatile.  Iodine  solutions,  therefore, 
should  be  kept  in  a  cool,  dark  place,  and  should  be  restandardized  at  suitable 
intervals. 


VOLUMETRIC  ANALYSIS  153 

2.  Standard  thiosulphate  solutions  may  also  be  prepared  by  dissolving 
the  theoretical  quantity  of  the  purified  salt  in  cold  water  (free  from  carbonic 
acid)  and  diluting  to  the  required  volume.  Such  solutions  are  quite  stable, 
and,  if  protected  from  carbon  dioxide,  in  a  cool,  dark  place,  they  may  be  kept 
for  months  without  appreciable  change. 

3.  Carbonic  acid  causes  a  slow  decomposition  of  the  thiosulphate  solu- 
tion, with  the  formation  of  free  sulphur  and  sulphurous  acid;  and,  since 
sulphurous  acid  acts  in  the  same  way,  the  decomposition  once  started 
becomes  progressive : 

Na&03+H.C03<^{N-2>;+^    H.SO,+S, 

and  Na,S,03+H.S03:!;:(N^^;+^    H.SO3+S. 

The  reducing  value  of  the  solution  increases  gradually  as  the  decomposition 
progresses ;  i.e.  the  solution  apparently  becomes  stronger.  When  it  is  con- 
sidered that  in  this  decomposition  each  molecule  of  thiosulphate  yields  one 
molecule  of  sulphite,  the  greater  reducing  value  is  readily  understood ;  for 

2Na2S203+l2= Na2S406+2NaI, 
while  iNa2S03+l2+H20=Na2S04+2HI. 

4.  Although  the  standardization  of  thiosulphate  solution  against  a 
weighed  quantity  of  pure  iodine  is  sometimes  advocated,  this  procedure  is 
by  far  too  troublesome ;  it  is  much  better  to* use  a  carefully  standardized 
iodine  solution.  Instead  of  this,  however,  a  definite  weight  of  potassium 
bromate,  iodate,  or  dichromate,  all  of  which  are  readily  obtainable  in  a 
pure  condition,  may  be  added  to  an  excess  of  pure  potassium  iodide  in 
slightly  acid  solution,  with  the  liberation  of  a  definite  quantity  of  iodine; 
and  this  may  be  used  in  the  standardization.  In  the  same  way,  the  thio- 
sulphate solution  may  also  be  standardized  indirectly  against  a  standard 
solution  of  potassium  permanganate. 

5.  Arsenious  oxide  dissolves  most  readily  in  caustic  alkalies,  and  for 
this  reason  the  sodium  hydroxide  is  used.  The  sodium  hydroxide,  however, 
must  subsequently  be  removed,  since  otherwise  it  would  react  with  the 
iodine ;  hence  the  acidification  with  hydrochloric  acid.  The  purpose  of  the 
bicarbonate,  which  under  the  analytical  conditions  is  without  action  upon 
the. iodine,  is  to  destroy  the  hydriodic  acid  formed  in  the  reversible  action, 
As203-f^  2  I2+  2  H2O  "^  As20s4-4  HI,  and  thus  drive  the  action  to  completion 
to  the  right.    The  reaction  may  then  be  written : 

Na2HAs03+l2+2  NaHC03=Na2HAs04+2  NaI4-2  CO2+H2O. 


154  QUANTITATIVE  CHEMICAL  ANALYSIS 

6.  Since  the  addition  of  iodine  in  excess  to  the  weakly  basic  bicarbonate 
solution  is  likely  to  lead  to  a  slight  degree  of  action,  it  is  best  in  this  titra- 
tion not  to  overstep  the  end-point. 

7.  Iodine  is  a  rather  expensive  chemical  and  it  is  well  worth  while  to 
recover  it  from  the  united  residues  of  a  large  class. 


VOLUMETRIC  ANALYSIS  I55 


THE  DETERMINATION  OF  ANTIMONY  IN  STIBNITE 

The  stibnite  should  be  an  ore  practically  free  from  arsenic 
and  iron,  and,  with  the  exception  of  a  siUceous  residue,  it  should 
be  wholly  soluble  in  hydrochloric  acid. 

Procedure.  Weigh  out  into  150  cc.  beakers  0.20  g.  portions 
of  the  finely  pulverized  ore,  add  to  each  about  0.3  g.  of  solid 
potassium  chloride  and  5  cc.  of  12-normal  hydrochloric  add 
(effervescence) J  and  allow  to  stand  for  10  minutes;  then  warm 
on  the  steam  bath  until  the  residue  is  white.  Add  2.5  g.  of  pul- 
verized tartaric  acid,  and  continue  the  heating  for  15  minutes; 
with  the  addition  of  a  little  of  the  hydrochloric  acid,  if  neces- 
sary, to  prevent  the  exposure  of  the  bottom  of  the  beaker. 

Now  add  water  to  the  mixture,  a  Httle  at  a  time,  with  care  to 
stop  if  there  is  any  indication  of  the  precipitation  of  antimony 
sulphide;  in  such  a  case,  warm  the  mixture  until  the  orange- 
red  substance  disappears,  and  again  proceed  with  the  gradual 
addition  of  water,  until  the  clear  liquid  has  a  volume  of  100  cc. 
Neutralize  the  cold  solution  with  sodium  hydroxide  (methyl 
orange),  and  then  add  6-normal  hydrochloric  acid,  2  drops  in 
excess. 

Transfer  the  solution  quantitatively  to  a  500  cc.  Erlenmeyer 
flask,  washing  out  the  beaker  with  three  25  cc.  portions  of 
sodium  bicarbonate  solution  (made  by  dissolving  20  g.  of  the 
pure  salt  in  300  cc.  of  cold  water),  add  5  cc.  of  starch  solution, 
and  titrate  with  the  standard  iodine  solution,  with  care  not  to 
overstep  the  end-point.  Report  the  percentage  of  antimony 
found. 

Notes.  —  i.  Stibnite  is  an  important  material  in  the  manufacture  of 
primers  for  cartridges.  The  pure  substance  has  the  composition  represented 
by  Sb2S3,  but  the  commercial  mineral  may  also  contain  the  compounds 
Sb20S2  and  Sb204,  besides  iron,  arsenic,  lead,  siliceous  gangue,  etc.  Upon 
warming  the  ore  with  hydrochloric  acid,  the  antimony  is  extracted,  and 
hydrogen  sulphide  evolved :  Sb2S3+6  HC1=  2  SbCls+s  H2S.  The  student 
will  recall,  however,  that  this  reaction  is  reversible ;  if  the  antimony  is  to 


156  QUANTITATIVE  CHEMICAL  ANALYSIS 

be  kept  in  solution,  the  expulsion  of  hydrogen  sulphide  must  be  complete 
before  the  final  dilution. 

2.  Antimony  trichloride,  in  the  presence  of  strong  hydrochloric  acid, 
is  slightly  volatile ;  with  potassium  chloride,  however,  it  forms  a  less  volatile 
double  salt.  With  careful  heating,  especially  in  the  presence  of  potassium 
chloride,  no  error  need  be  feared  from  this  source. 

3.  Antimony  trichloride  and  water  interact  to  give  insoluble  antimony! 
chloride,  SbOCl,  but  this  with  tartaric  acid  forms  a  soluble  antimonyl  com- 
pound similar  to  tartar  emetic : 

SbOCl-fH2C4H406=H(SbO)C4H406+HCl. 

The  tartaric  acid,  therefore,  is  added  to  prevent  the  precipitation  of  basic 
compounds  of  antimony  in  the  subsequent  treatment  of  the  solution.  If  a 
white  precipitate  should  be  obtained,  however,  the  solution  should  be 
rejected. 

4.  The  reaction  between  the  iodine  and  the  antimonyl  tartrate  is  not 
so  simple,  but  for  purposes  of  calculation  it  is  accurately  expressed  by  the 
equation,  Sb2034-2  I2+2  H20^Sb205+4  HI.  The  purpose  of  the  bicar- 
bonate is  here  also  to  neutralize  the  hydriodic  acid  formed,  and  thereby 
drive  the  oxidation  to  completion.  The  sodium  hydroxide  previously  added 
neutralizes  most  of  the  acid  and  makes  it  easy  to  provide  for  the  presence 
of  a  known  quantity  of  sodium  bicarbonate  in  the  titration;  the  solution 
should  be  distinctly  acid  when  the  bicarbonate  is  added. 

5.  If  the  ore  to  be  analyzed  contains  more  than  traces  of  iron,  it  is  dis- 
solved in  hydrochloric  acid,  the  antimony  precipitated  with  hydrogen  sul- 
phide, and  the  washed  precipitate  redissolved  in  hydrochloric  acid  and 
determined  as  above.  In  case  arsenic  also  is  present,  a  somewhat  more 
complicated  treatment  is  necessary. 


VOLUMETRIC  ANALYSIS  157 

THE  DETERMINATION  OF  CHROMIUM  IN  CHROMITE 

Procedure.  Weigh  out  into  bright  iron  crucibles  two  por- 
tions of  the  finely  pulverized  ore,  each  sufficient  to  contain 
about  45  mg.  of  chromium.  Also  weigh  out  roughly  upon  watch 
glasses  two  4  g.  portions  of  pure  sodium  peroxide.  (Owing  to 
the  tendency  of  the  peroxide  to  absorb  moisture,  the  first  por- 
tion should  be  mixed  with  one  sample  before  the  second  portion 
is  removed  from  the  container.)  Transfer  about  3  g.  of  the 
peroxide  to  each  crucible,  and  mix  this  with  the  ore  by  means  of  a 
dry  rod ;  remove  any  adhering  particles  from  the  rod  by  stir- 
ring with  it  the  remaining  peroxide,  and  transfer  the  latter 
to  the  surface  of  the  mixture.  Gradually  heat  the  crucible 
upon  a  triangle,  with  a  small  movable  flame,  until  the  mixture 
melts ;  and  continue  to  heat  at  about  this  temperature  for  5  or 
6  minutes.     Then  allow  the  melt  to  cool. 

In  a  tall  250  cc.  beaker,  treat  the  crucible  and  its  contents 
with  cold  water,  with  great  care  to  avoid  loss  by  effervescence. 
As  soon  as  the  action  slows  down,  heat  the  mixture  gradually, 
finally  at  the  boiling  temperature,  until  there  is  no  further  evo- 
lution of  oxygen.  Take  out  and  wash  off  the  crucible,  add 
sufficient  6-normal  hydrochloric  acid  (calculated)  to  almost 
neutralize  the  liquid,  and  filter;  receive  the  filtrate  in  a  large 
Erlenmeyer  flask.  To  the  filtrate  and  washings,  which  should 
be  slightly  alkaline,  add  about  0.2  g.  of  sodium  peroxide,  boil 
for  several  minutes,  and  acidify  with  6-normal  hydrochloric 
acid,  adding  5  cc.  in  excess.  The  volume  at  this  point  should 
be  about  200  cc.  Mix  with  this  solution  1.5  g.  of  pure  potassium 
iodide,  and  titrate  at  once  with  the  standard  thiosulphate  solu- 
tion until  the  brown  (iodine)  color  becomes  faint ;  then  add  5.  cc. 
of  starch  solution,  and  continue  the  titration  cautiously  until 
the  solution  becomes  pale  green  (CrCU)  with  no  tinge  of  blue. 
The  end-point,  though  very  sharp,  may  easily  be  passed  through 
carelessness.     It  is  well  to  have  a  white  surface  under  the  flask. 

Report  the  percentage  of  chromium  in  the  ore. 


158  QUANTITATIVE  CHEMICAL  ANALYSIS 

Notes.  —  i.  Fused  sodium  peroxide  attacks  most  materials;  although 
it  attacks  iron  and  nickel,  crucibles  of  these  metals  may  nevertheless  be 
used  in  this  case,  but  great  care  should  be  taken  not  to  raise  the  temperature 
too  high. 

2.  Chromite  is  an  ore  consisting  largely  of  ferrous  chromite,  Fe(Cr02)2. 
Fused  sodium  peroxide  oxidizes  the  ferrous  iron  to  sodium  ferrate  and  the 
chromium  to  sodium  chromate : 

2  (FeO  .  Cr203)  +  io  Na202=4  Na2Cr04+2  Na2Fe04+4  Na20. 

Upon  the  addition  of  water  to  the  melt,  the  chromate  dissolves,  the  ferrate 
is  decomposed  into  sodium  hydroxide,  ferric  oxide,  and  oxygen,  and  the 
excess  of  the  peroxide  yields  sodium  hydroxide  and  oxygen ;  the  liquid  is 
boiled  in  order  to  complete  the  destruction  of  the  peroxide,  since  if  present 
it  would  lead  to  the  formation  of  perchromate  upon  acidification. 

The  alkaline  chromate  solution  is  always  subject  to  a  slight  partial 
reduction  when  filtered  through  paper ;  hence  the  addition  to  the  filtrate  of 
a  little  sodium  peroxide,  followed  by  renewed  boiling. 

3.  The  mixture  should  not  be  acidified  before  filtration,  because  in  that 
case  iron  would  enter  the  solution  and  cause  large  errors  in  the  titration. 
The  partial  neutralization  is  to  prevent  the  destruction  of  the  filter  paper 
by  the  alkaline  liquid. 

4.  Upon  the  addition  of  potassium  iodide  in  excess  to  the  acidified 
fusion  extract,  iodine  is  liberated  quantitatively  according  to  the  equation, 

CrjOy-  -+6 1-+14  H+=  2  Cr+++H-7  H2O+3  I2. 

5.  Instead  of  employing  this  method,  it  is  of  course  possible  to  prepare 
the  acid  fusion  extract  as  described  in  the  procedure,  and  to  add  to  it  (in 
excess)  a  known  weight  of  pure  ferrous  ammonium  sulphate ;  the  excess  of 
this  substance  can  then  be  titrated  with  standard  dichromate  solution.  In 
that  case,  a  larger  sample  of  the  ore  may  be  taken  for  analysis. 


VOLUMETRIC  ANALYSIS  159 

THE  DETERMINATION  OF  LEAD  IN  AN  ORE 

Procedure.  Weigh  out  two  samples  of  the  finely  ground  ore 
sufficient  to  contain  about  0.20  g.  of  lead  (0.28-0.29  g.  of  a  70% 
ore),  and  treat  each  as  follows :  Moisten  the  sample  with  water, 
add  15  cc.  of  i2-normal  hydrochloric  acid,  and  evaporate  on 
the  steam  bath  to  about  5  cc.  Add  3  cc.  of  strong  nitric  acid, 
evaporate  nearly  to  dryness,  then  add  20  cc.  of  6-normal  hydro- 
chloric acid  and  again  heat  to  bring  all  the  lead  chloride  into 
solution.  Add  20  cc.  of  6-normal  sulphuric  acid  and  evaporate 
to  white  fumes.  Allow  to  cool,  add  50  cc.  of  water,  boil,  and 
then  add  15  cc.  of  alcohol ;  stir,  allow  to  settle,  and  filter.  Wash 
the  lead  sulphate  and  gangue  six  times  with  lo-cc.  portions  of 
0.5-normal  sulphuric  acid  (15  cc.  of  6-normal  acid  in  165  cc.  of 
water),  transfer  the  residue  to  a  small  beaker  by  means  of  a  jet 
of  water,  and  heat  it  gently  for  a  few  minutes  with  20  cc.  of 
ammonium  acetate  solution ;  ^  filter  the  liquid  through  the  orig- 
inal filter  and  wash  the  latter  with  small  portions  of  the  hot 
ammonium  acetate  solution.  Dilute  the  extract  to  150  cc, 
heat  to  boiling  and  add  from  a  pipette  10  cc.  of  potassium  dichro- 
mate  solution.^  Boil  the  mixture  gently  for  10  minutes,  filter 
off  the  precipitate  of  lead  chromate  and  wash  the  filter  and 
precipitate  about  ten  times  with  10  cc.  portions  of  dilute  am- 
monium acetate  solution  (25  cc.  of  the  extraction  solution  diluted 
to  250  cc),  until  the  excess  of  potassium  chromate  is  completely 
removed. 

Now  place  a  clean  500-cc.  Erlenmeyer  flask  under  the  funnel, 
and  with  a  jet  of  cold,  acid,  sodium  chloride  solution^  stir  up 
and  dissolve  the  precipitate;  continue  washing  with  the  same 
liquid  until  every  trace  of  color  is  removed  from  the  filter.  In 
any  case,  use  at  least  50  cc.  of  the  liquid.    Finally  dilute  to 

*  Made  by  neutralizing  30%  acetic  acid  with  6-normal  ammonia,  and  then  add- 
ing a  slight  excess  of  ammonia. 

*  A  solution  containing  75  g.  of  K2Cr207  per  liter. 

»  Mix  10  cc.  of  i2-normal  hydrochloric  acid  with  15  cc.  of  water,  and  add  this 
mixture  to  100  cc.  of  a  saturated  solution  of  sodium  chloride. 


i6o  QUANTITATIVE  CHEMICAL  ANALYSIS 

150  cc,  add  I  g.  of  potassium  iodide,  mix,  and  titrate  at  once 
with  a  solution  of  sodium  thiosulphate  (which  has  been  stand- 
ardized in  the  same  way  against  test  lead,  see  Note  5)  until 
the  brown  color  becomes  faint ;  then  add  5  cc.  of  starch  solution, 
and  continue  the  titration  cautiously  until  the  solution  becomes 
pale  green  (CrCla)  with  no  tinge  of  blue.  The  end-point  is  very 
sharp,  but  without  great  care  it  may  easily  be  passed.  It  is 
best  to  have  a  white  surface  under  the  flask. 
Report  the  percentage  of  lead  in  the  ore. 

Notes.  —  i.  The  ore  is  first  heated  with  strong  hydrochloric  acid  in 
order  to  expel  most  of  the  sulphur.  Nitro-hydrochloric  acid  is  then  used 
to  decompose  any  refractory  sulphides.  Upon  evaporating  the  chloride 
solution  to  white  fumes  with  sulphuric  acid,  the  volatile  acids  in  which  lead 
sulphate  is  slightly  soluble  are  completely  expelled,  and  upon  dilution  with 
water,  especially  if  alcohol  is  added,  the  lead  is  all  left  in  the  residue  as  lead 
sulphate. 

2.  Lead  sulphate  is  readily  dissolved  by  ammonium  acetate  solution, 
owing  to  the  exceptional  behavior  of  lead  acetate  with  respect  to  ionization 
(see  Part  I),  leaving  the  siliceous  gangue,  BaS04,  etc.,  as  a  residue. 

3.  While  lead  is  not  precipitated  from  solutions  containing  a  large  ex- 
cess of  acetate  ion  by  sulphates,  the  addition  of  a  soluble  chromate  causes 
the  precipitation  of  lead  chromate.  This  behavior  is  due  to  the  fact 
that  such  solutions  contain  Pb++-ion  at  an  extremely  low  concentration 
(owing  to  the  presence  of  the  lead  mainly  in  the  form  of  intermediate  or  com- 
plex ions,  as  (Pb.C2H302)+,  [Pb(C2H302)3l~,  etc.),  and  also  to  the  fact  that 
lead  sulphate  is  very  much  more  soluble  than  lead  chromate ;  the  lead-ion 
concentration  is  still  great  enough  in  such  solutions  to  cause  the  solubility 
product  of  lead  chromate  to  be  exceeded  upon  the  addition  of  potassium 
chromate  in  excess. 

4.  The  solubility  of  lead  chromate  in  the  acid  chloride  solution  is 
due  on  the  one  hand  to  the  lowered  concentration  of  the  chromate  ion, 
owing  to  the  formation  of  non-ionized  H2Cr04,  HCr04~,  etc.,  and  on  the 
other  hand  to  the  great  tendency  of  lead  ion  to  form  soluble  complexes 
with  chloride  solutions.  (Cf.  the  solubility  of  silver  chloride  in  chloride 
solutions.) 

5.  The  reaction  of  the  acid  solution  with  potassium  iodide  is  most  simply 
represented  by  the  equation, 

Cr207-+6 1-+14  H+=  2  Cr++++7  HaO-f  3  I2; 


VOLUMETRIC  ANALYSIS  i6i 

from  this  it  may  be  seen  that  one  atom  of  lead  (as  PbCr04)  leads  to  the 
liberation  of  3  atoms  of  iodine.  But,  since  the  composition  of  the  lead 
chromate  varies  slightly  with  the  conditions,  the  thiosulphate  solution  must 
be  standardized  under  identical  conditions  against  a  known  amount  of 
lead.  In  this  case,  0.20  g.  of  test  lead  should  be  dissolved  in  5  cc.  of 
6-normal  nitric  acid,  the  solution  evaporated  to  white  fumes  with  20  cc. 
of  6-normal  sulphuric  acid,  and  the  subsequent  operations  carried  out  as 
described  in  the  procedure. 


l62  QUANTITATIVE  CHEMICAL  ANALYSIS 

THE  DETERMINATION  OF  COPPER  IN  AN  ORE 

Principle.  This  method  is  based  upon  the  reaction  which 
takes  place  upon  the  addition  of  potassium  iodide  to  a  slightly 
acid  copper  salt  solution;  cuprous  iodide  is  precipitated  as  a 
cream-colored  powder,  and  iodine  is  set  free : 

2  CUSO4+4  KI=2  K2SO4  +  CU2I2+I2. 

The  iodine  is  promptly  titrated  with  a  standard  thiosulphate 
solution. 

Standardization  of  the  Thiosulphate  Solution.  Weigh  ac- 
curately two  portions  of  pure  bright  copper  wire  or  foil,  of  0.15- 
0.16  g.  each,  and,  in  250  cc.  Erlenmeyer  flasks,  dissolve  these  in 
5  cc.  portions  of  6-normal  nitric  acid.  Dilute  each  solution  to 
15  cc.  and  boil  to  expel  the  red  fumes ;  then  dilute  to  25  cc.  and 
add  ammonia  (sp.  gr.,  0.90)  in  slight  excess.  Again  boil  until 
the  ammonia  odor  is  faint,  add  80%  acetic  acid,  2-3  cc.  in  excess, 
and  boil  for  a  moment  longer,  agitating  the  flask  in  a  holder  to 
prevent  bumping.  Cool  to  room  temperature,  dilute  to  40  cc, 
add  a  solution  of  3  g.  of  potassium  iodide  in  10  cc.  of  water,  and 
titrate  at  once  with  the  approximately  tenth-normal  thiosul- 
phate solution  to  a  faint  brown  tinge ;  add  5  cc.  of  starch  solu- 
tion, and  continue  the  titration  until  the  last  faint  Hlac  tint  is 
removed  by  a  single  drop.  Do  not  overstep  the  end-point. 
From  the  data  obtained,  calculate  the  value  of  the  solution  per 
cubic  centimeter  in  terms  of  copper. 

Analytical  Procedure.  Weigh  out  into  300  cc.  beakers  samples 
of  the  ore  sufficient  to  furnish  about  0.15  g.  of  copper,  and  treat 
each  as  follows:  Add  10  cc.  of  hydrochloric  acid  (sp.  gr.,  1.19) 
and  5  cc.  of  nitric  acid  (sp.  gr.,  1.42)  and  heat  in  the  covered 
beaker  on  the  hot  plate  until  decomposition  is  complete,  adding 
more  of  the  acids  if  necessary,  and  enough  water  at  the  end  to 
hold  all  soluble  salts  in  solution.  Then  add  15  cc.  of  6-normal 
sulphuric  acid,  and  continue  the  heating  until  abundant  white 
fumes  begin  to  come  off.  _Cool,  add  50-60  cc.  of  water,  boil  for 


VOLUMETRIC  ANALYSIS  163 

a  moment,  and  allow  to  stand,  hot,  until  any  anhydrous  ferric 
sulphate  has  dissolved.  Finally,  filter  off  from  any  lead  sulphate, 
gangue,  and  sulphur,  receiving  the  filtrate  and  washings  in  a 
300  cc.  beaker.  Now  add  a  solution  of  5  g.  of  sodium  thio- 
sulphate  in  25  cc.  of  water,  boil  to  coagulate  the  precipitate, 
and  filter,  transferring  the  precipitate  quantitatively  to  the 
filter  by  means  of  hot  water.     Dry  the  precipitate  on  the  filter. 

Place  the  precipitate,  together  with  the  filter,  in  a  porcelain 
crucible,  ignite  gently  until  the  filter  is  consumed,  and  allow  to 
cool.  Transfer  the  bulk  of  the  precipitate  to  a  250  cc.  Erlenmeyer 
flask,  and  set  aside.  To  dissolve  the  last  portions  of  the  pre- 
cipitate from  the  crucible,  add  3  cc.  of  concentrated  nitric  acid 
and  2  cc.  of  water,  and  warm  gently  on  the  hot  plate,  finally 
pouring  the  acid  solution  into  the  flask  containing  the  bulk  of 
the  precipitate,  and  washing  out  the  crucible  with  a  few  small 
portions  of  6-normal  nitric  acid.  Heat  the  mixture  in  the  flask 
until  the  decomposition  is  complete,  dilute  to  25  cc,  boil,  add 
ammonia  in  slight  excess,  and  heat  until  the  odor  is  faint.  Add 
80%  acetic  acid,  2-3  cc.  in  excess,  and  boil  for  a  moment,  vigor- 
ously agitating  the  flask  to  prevent  bumping.  Cool  to  room 
temperature,  dilute  to  40  cc,  add  3  g.  of  potassium  iodide  dis- 
solved in  10  cc.  of  water,  and  titrate  at  once  with  the  thio- 
sulphate  solution,  as  previously  described.  Report  the  per- 
centage of  copper  in  the  ore. 

Notes.  —  i.  Since  iron  and  other  elements  likely  to  be  present  inter- 
fere with  the  process,  the  copper  must  be  separated  from  these.  Lead  is 
first  removed  by  means  of  sulphuric  acid,  after  which  the  copper  is  pre- 
cipitated from  the  hot,  acid  solution  by  means  of  sodium  thiosulphate ; 
this  gives  a  flocculent  precipitate  of  cuprous  sulphide  mixed  with  sulphur, 
which  filters  readily  and  can  be  washed  with  hot  water  without  fear  of 
oxidation.  Arsenic  and  antimony,  if  present,  are  also  precipitated,  but 
under  the  treatment  prescribed  the  usual  quantities  of  these  elements  are 
without  influence.  They  are  mostly  volatilized  during  the  ignition.  If 
antimony  is  present  in  appreciable  quantity,  it  is  perhaps  better  to  filter 
the  solution  before  the  addition  of  the  ammonia. 

2.  In  order  to  obtain  the  best  results  it  is  necessary  to  standardize  the 


1 64  QUANTITATIVE  CHEMICAL  ANALYSIS 

thiosulphate  solution  against  pure  metallic  copper.  When  this  is  done  the 
method  is  very  accurate ;  otherwise  the  results  are  not  so  good.  For  ex- 
ample, a  thiosulphate  solution  which  (titrated  against  a  freshly  stand- 
ardized iodine  solution)  had  a  calculated  copper  value  of  0.00608  g.  per 
cubic  centimeter,  was  found  upon  standardization  against  pure  copper  to 
have  a  value  of  0.0061 1  g.  per  cubic  centimeter. 

3.  Since  nitrous  fumes  liberate  iodine  from  potassium  iodide,  they  must 
be  completely  expelled  by  boiling  before  the  addition  of  the  salt.  The 
expulsion  of  the  last  traces  of  these  fumes  is  insured  by  boiling  the  solu- 
tion after  it  has  been  acidified  with  acetic  acid. 

4.  The  return  of  the  blue  tinge  in  the  liquid  after  long  standing  is  of  no 
significance,  but  a  quick  return  which  is  not  prevented  from  recurring  by 
the  addition  of  a  single  drop  of  the  thiosulphate  solution  is  usually  an 
evidence  of  faulty  work. 

5.  In  such  a  case,  or  if  the  end-point  has  accidentally  been  passed,  the 
same  sample  may  be  prepared  anew  for  titration :  Add  10  cc.  of  concen- 
trated nitric  acid,  and  heat  very  cautiously,  with  great  care  not  to  allow 
the  mixture  to  foam  over.  After  most  of  the  iodine  has  been  expelled, 
manipulate  the  flask  (in  a  holder)  over  a  free  flame  and  boil  the  solution 
down  rapidly  to  a  volume  of  5-10  cc.  Dilute  to  25  cc.  with  water,  boil, 
add  ammonia  in  slight  excess,  and  finish  as  described  in  the  procedure. 

6.  In  the  electrolytic  determination  of  copper  in  ores  containing  arsenic 
and  other  interfering  substances,  a  satisfactory  copper  solution  is  most 
readily  prepared  by  dissolving  the  ignited  thiosulphate  precipitate  in  a 
suitable  quantity  of  strong  nitric  acid,  with  subsequent  dilution  to  the 
required  volume. 


VOLUMETRIC  ANALYSIS  165 

C.   PRECIPITATION  METHODS 

General  Discussion.  Perhaps  the  simplest  example  of  a  pre- 
cipitation process  is  that  furnished  by  the  method  for  silver 
which  was  originated  by  Gay-Lussac  in  1832,  and  which  is  still 
widely  used  in  determining  the  fineness  of  silver  bullion.  This 
process  is  based  upon  the  reaction  between  silver  nitrate  and  a 
standard  solution  of  sodium  chloride,  which  runs  to  completion 
in  consequence  of  the  union  of  silver-ion  with  chloride-ion  to 
furnish  insoluble  silver  chloride. 

When  the  silver  chloride  first  separates  it  is  finely  divided, 
and  a  very  minute  quantity  can  easily  be  recognized.  Upon 
vigorously  shaking  the  solution,  the  precipitate  coagulates  and 
settles,  leaving  the  supernatant  liquid  bright  and  clear.  Hence, 
if  silver  nitrate  is  titrated  with  a  solution  of  sodium  chloride, 
with  shaking  in  a  glass-stoppered  bottle  after  each  addition, 
the  point  at  which  the  further  addition  of  the  standard  solution 
ceases  to  produce  a  precipitate  can  readily  be  determined. 
Near  the  end-point  it  is  customary  to  use  a  standard  solution 
of  one  tenth  the  value  of  that  used  at  the  start. 

In  the  case  of  this  reaction,  this  method  of  determining  the 
end-point  admits  of  a  very  high  degree  of  accuracy,  and  it  is 
the  method  in  use  at  the  government  mints.  Since,  however, 
it  is  rather  tedious  and  demands  considerable  skill  and  experi- 
ence, a  slightly  less  accurate  but  much  more  convenient  method 
is  generally  employed. 

Silver  thiocyanate  is  perfectly  white,  and  even  less  soluble 
than  silver  chloride;  it  is  therefore  possible  to  titrate  silver 
very  accurately  with  a  standard  solution  of  an  alkali  thiocyanate. 
If  the  solution  contains  ferric  alum,  and  also  nitric  acid  to  pre- 
vent its  hydrolysis,  the  addition  of  the  slightest  excess  of  the 
thiocyanate  solution  can  be  readily  recognized  by  the  pink 
tint  which  it  imparts  to  the  mixture.  This  method  (Volhard^s) 
is  also  suitable  for  the  determination  of  the  halogens  (except 
fluorine)  and  of  certain  other  ions  which  give  silver  compounds 


1 66  QUANTITATIVE  CHEMICAL  ANALYSIS 

insoluble  in  dilute  nitric  acid.  A  measured  volume  of  standard 
silver  nitrate  solution  is  added  in  excess,  and  the  excess  subse- 
quently determined  by  means  of  the  standard  thiocyanate 
solution. 

THE  PREPARATION  AND  STANDARDIZATION  OF  APPROXI- 
MATELY ONE  TENTH  NORMAL  SOLUTIONS  OF  SILVER 
NITRATE  AND  AMMONIUM  THIOCYANATE 

Procedure.  Dissolve  4.0  g.  of  ammonium  thiocyanate  (or 
5.0  g.  of  the  potassium  salt)  in  water  and  dilute  the  solution  to 
500  cc.  Also  dissolve  8.5  g.  of  silver  nitrate  in  water  and  dilute 
the  solution  to  500  cc.  Further,  mix  10  cc.  of  6-normal  nitric 
acid  with  40  cc.  of  water,  heat  the  solution  to  boiling,  and  dis- 
solve in  the  hot  liquid  5  g.  of  pure  ferric  alum.  Allow  the  solu- 
tion to  cool,  and  keep  it  for  use  as  an  indicator. 

Now  fill  the  burettes  with  the  respective  solutions,  placing 
the  silver  nitrate  solution  in  a  glass-stoppered  burette.  (Ob- 
serve the  usual  precautions,  and  place  all  solutions  and  precipi- 
tates containing  silver  in  the  receptacle  for  silver  residues,)  Run 
out  20  cc.  of  the  silver  nitrate  into  an  Erlenmeyer  flask,  dilute 
to  150  cc,  add  10  cc.  of  freshly  boiled  6-normal  nitric  acid,  and 
5  cc.  of  the  indicator.  With  vigorous  shaking,  run  in  the  thio- 
cyanate solution  until  a  faint  pink  tinge  is  imparted  to  the  mix- 
ture. If  the  end-point  is  overstepped,  titrate  back  with  the  silver 
nitrate  solution.  Calculate  the  ratio  of  the  thiocyanate  to  the 
silver  nitrate  solution,  and  repeat  the  operation  until  the  values 
found  agree  within  two  parts  in  one  thousand. 

Finally,  standardize  the  silver  nitrate  solution,  as  follows: 
Weigh  out  portions  of  pure  sodium  chloride,  of  0.12-0.14  g. 
each,  dissolve  these  in  75-cc.  portions  of  water,  heat  to  boihng, 
and  with  stirring  run  into  each  from  a  burette  25.00  cc.  of  the 
silver  nitrate  solution.  Add  10  cc.  of  freshly  boiled  6-normal 
nitric  acid,  stir,  and  filter,  washing  the  precipitate  by  decantation 
with  several  small  portions  of  hot  distilled  water,  and  pouring 
these  slowly  over  the  filter;    the  united  filtrate  and  washings 


VOLUMETRIC  ANALYSIS  167 

should  have  a  volume  of  about  150  cc.  To  this  solution  add 
5  cc.  of  the  indicator,  and  titrate  the  excess  of  silver  with  the 
thiocyanate  solution,  as  already  described.  From  the  data 
obtained,  calculate  the  normaUty  factor  of  the  silver  nitrate 
solution;  and  from  the  mean  of  the  duphcate  values,  which 
should  agree  within  two  parts  in  a  thousand,  calculate  the  nor- 
maUty factor  of  the  thiocyanate  solution. 

Notes.  —  i.  The  reactions  between  the  thiocyanate  and  the  indicator 
are  essentially  as  follows : 

Fe++++6  CNS-  :$  Fe(CNS)3+3  CNS"  t^  [Fe(CNS)6] — . 

It  will  be  recalled  that  in  testing  for  ferric  iron  with  potassium  thiocyanate, 
it  is  necessary  to  add  a  large  excess  of  the  latter  in  order  to  detect  the  smallest 
possible  quantity  of  iron.  In  the  same  way,  when  using  ferric  iron  as  an 
indicator  for  thiocyanate,  it  is  necessary  to  provide  a  high  concentration 
of  the  former  in  order  to  detect  the  slightest  possible  excess  of  the  thio- 
cyanate in  the  solution.  The  reactions  which  give  rise  to  the  colored  sub- 
stances are  reversible,  but  in  the  presence  of  a  large  excess  of  one  of  the 
colorless  constituents  the  dissociation  of  the  colored  substances  is  prevented 
by  mass  action. 

2.  Nitric  acid  is  added  to  the  solution  to  be  titrated  in  order  to  prevent 
the  hydrolysis  of  the  ferric  salt,  which  would  impart  a  brownish ired  color 
to  the  mixturfe.  ^  It  is  bo^d  to  free  it  from  nitrous  fumes,  though  this  is  of 
less  importance  nere  thaWjy^ting  for  iron  in  the  presence  of  nitric  acid ; 
nitrous  fumes  color  the  thiOTjKiate  pink. 

3.  Sodium  chloride  may  easily  be  obtained  pure  by  filtering  a  concen- 
trated solution  of  the  commercial  salt,  saturating  it  with  hydrogen  chloride 
gas,  and  filtering  oflf  the  precipitate.  The  latter  is  washed  with  strong 
hydrochloric  acid  and  dried  at  150°,  or  higher.  -  • 

4.  Standard  solutions  of  silver  nitrate  can  of  course  be  prepared  by  the 
solution  of  the  calculated  amount  of  pure  metaUic  silver  in  nitric  acid,  and 
dilution  to  the  required  volume;  or  by  means  of  the  calculated  weight  of 
pure  silver  nitrate. 


i68  QUANTITATIVE  CHEMICAL  ANALYSIS 


THE  DETERMINATION  OF  CHLORINE  IN  A 
SOLUBLE  CHLORIDE 

The  sample  may  be  an  artificial  mixture  of  the  chloride  and 
carbonate  of  sodium. 

Procedure.  Weigh  out  into  300-cc.  beakers,  two  portions, 
each  sufficient  to  contain  about  0.12  g.  of  sodium  chloride,  and 
treat  each  as  follows:  Dissolve  the  sample  in  50  cc.  of  water, 
run  in  from  a  burette  30.00  cc.  of  the  standard  silver  nitrate 
solution,  and  carefully  acidify  the  mixture  with  dilute  nitric 
acid.  Heat  to  boiling,  see  that  the  Hquid  is  distinctly  acid,  and 
filter.  Receive  the  filtrate  and  washings  in  a  300-cc.  Erlenmeyer 
flask.  To  the  united  filtrate  and  washings,  which  should  have 
a  volume  of  about  150  cc,  add  10  cc.  of  6-normal  nitric  acid  and 
5  cc.  of  the  indicator  solution,  and  titrate  with  the  standard 
thiocyanate  solution,  as  already  described.  Calculate  the  per- 
centage of  chlorine  in  the  sample. 

Notes.  —  i.  Since  silver  chloride  is  several  times  as  soluble  as  silver 
thiocyanate,  the  former  must  be  filtered  off  before  the  titration  of  the 
excess  of  silver  nitrate ;  otherwise  the  silver  chloride  would  react  with  the 
thiocyanate  solution  and  render  the  end-point  uncertain.  This  behavior 
is  best  represented  by  the  following  system  of  equilibria : 

AgCl^AgCi:^^      /C1-+ 

(Solid)  (Diss'd)        \Ag+     (+CNS-^AgCNS:!^AgCNS). 

{Diss'd)         (Solid) 

That  is,  if  the  silver  chloride  were  left  in  the  mixture  during  the  titration, 
owing  to  the  slow  conversion  of  the  soluble  (colored)  thiocyanate  com- 
pounds into  insoluble  silver  thiocyanate,  there  would  be  no  permanent 
end-point. 

2.  Silver  bromide  and  silver  iodide  are  less  soluble  than  silver  thio- 
cyanate, so  that  in  the  determination  of  bromine  and  iodine  by  this  method 
it  is  not  necessary  to  filter. 

3.  Soluble  chlorates,  etc.,  may  be  determined  by  this  method  by  first 
reducing  them  to  the  corresponding  halides  {e.g.  with  sulphurous  acid), 
and  then  determining  the  latter. 


VOLUMETRIC  ANALYSIS  169 

4.  For  a  simple  rapid  method  for  the  determination  of  halogen  in  organic 
substances,  with  the  final  use  of  Volhard's  procedure,  see  W.  A.  Van  Winkle 
and  G.  McP.  Smith,  Jour.  Amer.  Chem.  Soc,  vol.  42,  pp.  333-47  (1920). 

5.  For  other  uses  of  precipitation  methods,  see  Part  V,  Problems  90, 
91,  92,  and  93. 


PART    IV 

QUESTIONS 

Exercises  with  the  Balance. 

1.  What  is  the  purpose  of  weighing? 

2.  Explain  the  mechanical  theory  of  the  balatice. 

3.  Give  five  conditions  which  must  be  satisfied  by  a  good  balance. 

4.  What  conditions  must  be  fulfilled  in  order  that  a  balance  may  be 
considered  properly  adjusted  for  use  ? 

5.  How  may  the  rest-point  of  a  balance  be  determined  ?    Illustrate. 

6.  Explain  the  following  methods  of  weighing:  {a)  ordinary  method; 
(b)  weighing  by  the  use  of  deflections;  (c)  weighing  by  transposition; 
(d)  weighing  by  substitution. 

7.  Describe  a  procedure  for  the  calibration  of  a  set  of  weights. 

8.  Discuss  the  errors  in  weighing  which  may  be  due :  (a)  to  inequalities 
in  length  in  the  beam  arms ;   (b)  to  the  buoyancy  of  the  atmosphere. 

9.  What  is  a  desiccator?  Explain  why  it  is  necessary,  and  give  the 
principles  upon  which  its  use  is  based. 

The  Determination  of  Chlorine. 

1.  What  substances  would,  if  present,  interfere  with  this  determination  ? 

2.  Why  should  the  solution  be  acidified  with  nitric  acid  ?  Why  should 
a  large  excess  of  nitric  acid  be  avoided  ? 

3.  Why  should  the  solution  not  be  heated  until  after  the  addition  of  the 
silver  nitrate  ?    Why  is  it  then  heated  ? 

4.  What  are  the  advantages  of  washing  by  decantation?  In  washing  a 
precipitate,  whether  by  decantation  or  otherwise,  why  should  the  liquid  each 
time  be  removed  as  far  as  possible  before  the  addition  of  fresh  wash  hquid  ? 

5.  How  can  you  tell  when  the  precipitate  has  been  sufiiciently  washed  ? 

6.  Why  is  the  filter  paper  ignited  separately  from  the  bulk  of  the  pre- 
cipitate ?  What  is  the  object  of  the  treatment  with  nitric  and  with  hydro- 
chloric acid  ?    Explain. 

7.  What  is  the  efifect  of  Hght  upon  silver  chloride?  Is  the  action  of 
diffused  dayhght  a  serious  source  of  error? 

170 


QUESTIONS  171 

8.  Why  should  the  precipitate  be  heated  until  it  just  begins  to  fuse? 
What  is  the  effect  of  overheating?    Of  underheating ? 

9.  Water  saturated  with  silver  chloride  at  100°  contains  about  22  mg. 
of  the  salt  per  liter.  Explain  why  the  precipitate  may  be  thoroughly  washed 
with  hot  water  without  undue  loss.  What  is  the  solubihty  of  silver  chloride 
in  water  at  the  ordinary  room  temperature  ?  In  the  precipitation  of  chloride - 
ion,  why  should  silver  nitrate  be  added  in  moderate  excess? 

10.  Given  an  aqueous  solution  of  silver  chloride  in  equilibrium  with  a 
quantity  of  the  solid  salt :  What  would  happen  upon  the  addition  of  (a)  a 
few  drops  of  silver  nitrate  solution  ?  (6)  a  few  drops  of  sodium  chloride 
solution?     (c)  a  large  excess  of  sodium  chloride?    Explain  in  each  case. 

11.  How  may  the  crucible  safely  and  readily  be  cleaned  after  the  ignition 
of  the  silver  chloride  ?    Write  equations  to  show  the  reactions  involved. 

12.  (a)  Explain  the  solubiHty  of  silver  chloride  in  each  of  the  following 
substances :  aqueous  ammonia ;  potassium  cyanide  solution ;  sodium  thio- 
sulphate  solution,  (b)  Is  silver  iodide  soluble  in  aqueous  ammonia?  In 
a  solution  of  potassium  cyanide  ?    Explain  your  answers. 

13.  What  other  substances  may  be  determined  in  a  similar  manner  in 
the  form  of  insoluble  silver  salts  ? 

14.  Why  are  Gooch  crucibles  preferable  to  ordinary  paper  filters;  es- 
pecially, for  example,  in  the  determination  of  iodides,  cyanides,  etc.  ? 

15.  Starting  with  the  native  mineral,  describe  the  treatment  which 
renders  the  asbestos  suitable  for  use  in  the  preparation  of  Gooch  crucibles. 

The  Determination  of  Iron  and  of  Sulphur  in  a  Soluble  Sulphate 

of  Iron 

Iron.  —  I.  A  mixture  consists  of  ferric  sulphate,  sodium  carbonate,  and 
potassium  sulphate,  each  of  which  is  soluble  in  cold  water.  Will  the  mix- 
ture dissolve  in  water?    Illustrate  your  answer  by  means  of  equations. 

2.  How  is  the  sample  taken  into  solution  for  this  analysis? 

3.  In  the  precipitation  with  ammonium  hydroxide,  why  must  the  iron 
be  present  wholly  in  the  ferric  condition  ? 

4.  How  is  the  ferrous  iron  oxidized  in  this  analysis?  Write  the  equa- 
tion.   How  can  you  tell  when  the  oxidation  is  complete? 

5.  How  may  it  be  ascertained  whether  the  original  sample  contains 
ferrous  iron  ?    Is  it  worth  while  to  make  this  test  ?    Why  ? 

6.  Explain  in  full  the  principle  involved  in  the  use  of  double  precipita- 
tion for  effecting  the  complete  separation  of  sulphate-ion  from  the  ferric 
iron. 

7.  Why  is  the  filter  paper  macerated  with  the  solution  before  the  second 
precipitation  of  the  iron? 


172  QUANTITATIVE  CHEMICAL  ANALYSIS 

8.  Is  it  necessary  to  completely  wash  out  the  ammonium  chloride 
before  the  ferric  hydroxide  is  ignited ?    Why? 

9.  What  precautions  are  to  be  observed  in  igniting  the  precipitate  of 
ferric  hydroxide  ? 

10.  Name  two  other  metals  which  may  be  determined  by  a  similar  pro- 
cedure. What  additional  precautions  should  be  taken  in  their  determina- 
tion, and  why? 

11.  In  the  determination  of  these  three  metals,  would  it  be  equally  well 
to  use  sodium  hydroxide  as  the  precipitating  reagent  ?  Give  the  reason  for 
your  answer  in  the  case  of  each  metal. 

12.  What  is  the  effect  of  tartaric  or  citric  acid,  sugars,  etc.  upon  the  pre- 
cipitation of  ferric  hydroxide  by  means  of  ammonium  or  sodium  hydroxide  ? 
Do  any  other  metalHc  ions  behave  like  Fe''"+"*"-ion  in  this  respect  ?    Explain. 

Sulphur.  —  I.  Why  must  nitrates  be  removed  before  the  precipitation 
with  barium  chloride  ?    How  is  this  done  ?    Write  the  reaction. 

2.  Name  some  other  substances  which,  if  present,  should  be  removed 
before  precipitation  with  barium  chloride. 

3.  Why  is  barium  chloride  chosen  as  the  reagent  for  sulphate,  rather 
than  lead  nitrate  or  strontium  chloride  ? 

4.  How  many  cubic  centimeters  of  i -normal  barium  chloride  solution 
will  be  required  to  precipitate  the  sulphate  from  one  gram  of  a  sample  con- 
taining 60%  of  Fe2(S04)3  and  14%  of  K2SO4? 

5.  What  are  the  correct  conditions  for  the  precipitation  of  the  sulphate? 
Explain  in  full. 

6.  What  precautions  should  be  observed  in  the  ignition  of  the  barium 
sulphate,  and  why? 

7.  What  other  substances  may  be  determined  as  insoluble  sulphates? 
What  reagent  is  used  in  their  precipitation? 

The  Determination  of  Sulphur  in  a  Sulphide  Ore. 

1.  Assuming  complete  oxidation,  write  an  equation  to  show  the  action 
of  nitric  acid  upon  iron  pyrites,  FeS2. 

2.  If  pure  FeS2  were  decomposed  in  this  way  with  nitric  acid,  the  iron 
precipitated  with  ammonia,  and  the  filtrate  evaporated  with  hydrochloric 
acid  on  the  hot  plate  to  dryness  (to  remove  the  nitrates),  would  there  be  any 
danger  of  losing  a  portion  of  the  sulphur?    Explain. 

3.  Explain  the  solubility  of  lead  chloride  in  ammonium  chloride  solution. 
(Cf .  the  behavior  of  silver  chloride.) 

4.  In  the  analysis  of  an  ore  containing  lead,  how  may  we  prevent  the 
precipitation  of  a  portion  of  the  sulphur  as  lead  sulphate?  Fully  explain 
your  answer. 


QUESTIONS  173 

5.  Outline  an  experimental  procedure  for  the  determination  of  sulphur 
in  heavy  spar,  BaS04. 

The  Determination  of  Potash. 

1.  Is  this  a  precipitation  method?  If  not,  to  what  class  does  it 
belong? 

2.  Why  is  it  necessary  to  remove  hydrochloric  and  sulphuric  acids  before 
the  extraction  of  the  soluble  perchlorates ?  How  are  these  acids  removed? 
Explain. 

3.  What  weight  of  HCIO4  will  be  required  to  convert  one  gram  of  NaCl 
into  NaC104?  One  gram  of  KCl  into  KCIO4?  What  connection  exists 
between  these  questions  and  the  procedure  ?     (Cf .  Part  V,  Problem  30.) 

4.  Explain  the  fact  that  phosphates,  even  if  insoluble  in  alcohol,  need 
not  be  removed  before  making  the  extraction.  What  precaution  should  be 
observed  when  phosphates  are  present,  and  why  ? 

5.  Do  any  of  the  following  salts  interfere  with  the  method ;  and,  if  so, 
how  is  the  difficulty  overcome?    NH4CIO4,  Ba(C104)2,  Mg(C104)2. 

6.  Why  is  a  small  quantity  of  HCIO4  added  to  the  alcohol  to  be  used  in 
the  extraction?  If  50  cc.  of  this  liquid,  capable  of  dissolving  2  mg.  of  pure 
KCIO4,  are  used  in  the  extraction,  why  is  not  the  quantity  of  KCIO4  found 
in  the  residue  2  mg.  less  than  the  true  value  ? 

7.  How  may  we  determine  the  point  at  which  the  extraction  is  complete  ? 
Explain. 

The  Analysis  of  Limestone. 

1.  What  is  the  principal  component  of  limestone?  What  other  com- 
pounds are  usually  present  ?    Why  is  it  important  to  analyze  a  limestone  ? 

2.  Explain  briefly  the  method  described  in  the  procedure  for  the  deter- 
mination of  carbon  dioxide. 

3.  What  are  the  objections  to  this  method?  Under  what  conditions 
can  the  method  be  relied  upon  to  furnish  accurate  results  ? 

4.  What  modifications  should  be  made  in  the  procedure  if  the  sample 
to  be  analyzed  is  a  baking  powder? 

5.  Explain  the  operation  of  the  aspirator.  Why  is  the  carbon  dioxide 
more  readily  removed  when  air  is  bubbled  through  the  solution?  (Cf. 
Part  I,  The  Evaporation  of  Liquids.) 

6.  What  method  for  the  determination  of  carbon  dioxide  may  be  re- 
garded as  the  converse  of  the  method  described  in  the  procedure?  De- 
scribe it. 

7.  Can  you  think  of  any  other  method  that  might  serve  for  the  exact 
determination  of  carbon  dioxide?    (Cf.  Part  V,  Problem  33.) 


174  QUANTITATIVE  CHEMICAL  ANALYSIS 

8.  In  preparing  a  solution  of  the  limestone,  for  the  determination  of 
calcium  oxide,  etc.,  what  is  the  reason  for  the  double  evaporation  with 
hydrochloric  acid,  followed  by  continued  heating  on  the  steam  bath  ?  When 
this  residue  is  extracted  with  hydrochloric  acid,  what  is  the  insoluble  material 
that  is  left  ? 

9.  Why  is  the  insoluble  residue  first  washed  with  dilute  acid  rather 
than  with  water  ?    Would  cold  water  do  about  as  well  ?    Why  ? 

10.  What  is  the  purpose  of  adding  bromine  water  in  the  precipitation 
with  ammonium  hydroxide?  Why  must  a  large  excess  of  ammonia  be 
avoided  ? 

11.  Why  should  the  ammonium  hydroxide  precipitate  be  filtered  off 
promptly?  Of  what  does  this  precipitate  consist ?  How  is  this  precipitate 
treated,  and  why? 

12.  How  is  the  first  ammonium  oxalate  precipitate  treated,  and  why? 

13.  Explain  why  so  Httle  ammonium  oxalate  solution  is  added  in  the 
second  precipitation  of  the  calcium. 

14.  Explain  the  solubility  of  calcium  oxalate  in  hydrochloric  acid  and 
its  insolubility  in  acetic  acid. 

15.  What  reactions  take  place  when  calcium  oxalate,  CaC204 .  H2O, 
is  ignited? 

16.  Why  should  the  solution  be  only  faintly  ammoniacal  for  the  pre- 
cipitation of  the  magnesium?  Why  is  more  ammonia  later  added  to  the 
precipitation  mixture?  Why  is  the  precipitation  mixture  then  allowed 
to  stand  for  several  hours  before  filtration  ? 

17.  How  is  the  first  precipitate  of  magnesium  ammonium  phosphate 
treated?  Explain  why  this  is  necessary.  Is  it  possible  to  accomplish  the 
same  result  in  any  other  way,  and  if  so  how? 

18.  Why  is  the  precipitate  washed  with  dilute  ammonia  rather  than 
with  water? 

19.  Write  an  equation  to  show  what  happens  when  MgNH4P04 .  6  H2O 
is  ignited;  when  Mg[(NH4)2P04]2  •  n  H2O  is  ignited. 

20.  What  precautions  should  be  observed  in  igniting  the  precipitate, 
and  why  ? 

The  Determination  of  Phosphoric  Anhydride  in  Phosphate  Rock. 

1.  What  are  the  chief  components  of  phosphate  rock?  What  other 
compounds  are  usually  present  ?    What  is  apatite  ? 

2.  Explain  by  means  of  the  ionic  theory  and  the  solubiHty  product  law 
the  fact  that  calcium  phosphate,  Ca3(P04)2,  will  dissolve  in  nitric  acid. 

3.  Why  is  it  necessary  to  remove  any  soluble  silicic  acid  that  may  be 
present  in  the  nitric  acid  solution  of  the  mineral  ?    How  is  this  done  ? 


QUESTIONS  175 

4.  Explain  why  it  is  possible,  in  washing  the  insoluble  residue  of  silica, 
to  make  the  test  for  phosphate  in  the  washings  by  the  addition  of  ammonia 
alone. 

5.  Why  is  it  directed  to  neutrahze  the  nitric  acid  solution  with  ammonia 
and  then  to  make  it  slightly  acid  with  nitric  acid,  before  the  addition  of  the 
molybdate  reagent  ? 

6.  Why  is  an  acid  solution  of  ammonium  nitrate  used  in  washing  the 
yellow  precipitate?  (Cf.  question  9.)  What  would  happen  if  this  am- 
monium nitrate  wash-liquid  were  alkaline  with  ammonia  ?    Explain. 

7.  Could  the  phosphorus  be  determined  by  igniting  and  weighing  the 
yellow  precipitate  ? 

8.  Why  do  we  first  precipitate  the  phosphate  with  molybdate  solution 
instead  of  precipitating  it  directly  from  the  original  solution  with  magnesia 
mixture  ? 

9.  Explain  why  a  large  excess  of  the  molybdate  reagent  is  necessary 
for  the  complete  precipitation  of  the  phosphate. 

10.  Show  by  means  of  an  equation  the  reaction  between  ammonium 
phosphomolybdate,  (NH4)3P04.i2  M0O3,  and  ammonium  hydroxide. 
Sodium  hydroxide. 

11.  What  is  magnesia  mixture?  What  is  the  purpose  of  the  ammonium 
chloride  ?    Explain. 

12.  Why  should  the  magnesium  ammonium  phosphate  precipitation 
mixture  be  allowed  to  stand  for  some  hours  before  filtering? 

13.  Why  is  dilute  ammonia  used  in  washing  the  precipitate  of  mag- 
nesium ammonium  phosphate  ?    Why  not  use  water  ? 

14.  What  precautions  should  be  observed  in  the  ignition  of  this  pre- 
cipitate, and  why? 

15.  Name  another  substance  which  may  be  determined  by  precipitation 
with  magnesia  mixture.  Can  any  elements  be  determined  in  a  similar 
manner  by  precipitation  with  sodium  or  ammonium  phosphate?  If  so, 
name  them. 

16.  If  it  were  desired  to  determine  the  phosphorus  in  a  sample  of  steel, 
in  which  it  is  present  as  iron  phosphide,  what  would  be  the  procedure? 
Explain. 

The  Determination  of  Silica  in  a  Refractory  Silicate. 

1.  Why  is  it  essential  to  grind  the  whole  of  the  sample  very  fine? 

2.  Write  an  equation  to  show  the  reaction  between  orthoclase,  KAlSisOs, 
and  sodium  carbonate,  above  the  melting  point  of  the  latter.  Why  is  a 
very  large  excess  of  the  latter  used  ? 

3.  How  can  you  tell  whether  the  decomposition  is  complete  {a)  by  in- 


176  QUANTITATIVE  CHEMICAL  ANALYSIS 

specting  the  mixture  during  fusion;    {b)  after  the  treatment  of  the  melt 
with  dilute  acid  ? 

4.  Why  is  the  fused  mass  treated  with  a  considerable  volume  of  dilute 
acid,  rather  than  with  concentrated  acid  ? 

5.  How  is  the  silica  separated  from  the  residue  left  upon  evaporation 
with  dilute  hydrochloric  acid  ?    Explain. 

6.  What  is  the  purpose  of  treating  the  sihca  with  hydrofluoric  and  sul- 
phuric acids?    Explain  in  full. 

7.  How  would  you  determine  the  percentage  of  mixed  iron  and  aluminum 
oxides  in  an  insoluble  silicate?  How  would  you  determine  the  calcium 
oxide?    The  magnesium  oxide ? 

8.  How  would  you  determine  the  sihca  in  a  sihcate  mineral  which  may 
be  readily  decomposed  by  hydrochloric  acid? 

The  Electrolytic  Determination  of  Copper. 

1.  Define:  ampere;  volt;  ohm.  State  Ohm's  law.  State  Faraday's 
laws. 

2.  Show  by  means  of  a  diagram  how  the  apparatus  is  assembled  for 
an  electrolytic  determination. 

3.  Explain  the  electrolytic  reduction  of  nitric  acid  to  ammonia. 

4.  Explain  the  deposition  of  lead  peroxide  upon  the  anode,  in  the  elec- 
trolysis of  a  solution  containing  lead  nitrate  and  nitric  acid. 

5.  What  is  meant  by  the  discharge  potential  of  an  ion?  The  decom- 
position voltage  of  a  salt  ?    What  is  polarization  ?    Explain  in  each  case. 

6.  Explain  why  it  is  possible  to  separate  copper  from  nickel  by  elec- 
trolysis ?     Can  nickel  be  separated  from  cobalt  in  this  way  ?     Why  ? 

7.  What  are  the  advantages  of  mechanical  stirring  during  electrolysis? 

8.  What  is  meant  by  current  density?  What  is  the  unit  of  current 
density?  Why  is  current  density  a  factor  of  the  greatest  importance  in 
electro-analysis  ? 

9.  What  factors  favor  the  formation  of  a  satisfactory  deposit?  What 
factors  interfere  with  it  ? 

10.  Outline  a  method  for  the  preparation  of  a  solution  suitable  for 
electrolysis  when  the  sample  to  be  analyzed  is  a  copper  ore  containing 
arsenic.    Why  should  the  arsenic  be  removed  ? 

11.  Discuss  the  materials  from  which  electrodes  may  be  prepared,  and 
the  form  of  the  electrodes  and  electrolytic  vessels. 

12.  Name  the  factors  in  the  electrolytic  work  which  should  receive 
especial  attention  in  an  endeavor  to  make  accurate  and  reliable  copper 
determinations  as  rapidly  as  possible.  Explain  in  the  case  of  each  factor 
mentioned. 


QUESTIONS  177 " 

Volumetric  Analysis :  Fundamental  Principles. 

1.  Specify  the  chief  uses  of  measuring  flasks;  of  transfer  pipettes;  of 
burettes. 

2.  Describe  the  preparation  of  sulphuric  acid-dichromate  cleaning  solu- 
tion.   How  is  it  used  ? 

3.  In  a  titration  with  a  solution  correctly  made  up  to  tenth-normal  con- 
centration at  20°,  a  burette  correctly  graduated  for  use  at  20°  is  used  at  an 
actual  temperature  of  27.5°,  and  the  indicated  volume  of  solution  withdrawn 
is  27.68  cc. ;  to  how  many  cubic  centimeters  of  tenth-normal  solution  does 
this  liquid  correspond  ?     (See  Part  I.) 

4.  Define  the  term  "liter." 

5.  Describe  a  method  for  the  calibration  of  a  loo-cc.  measuring  flask. 
Illustrate  your  description.     (See  Part  I.) 

6.  What  is  a  standard  solution?  A  normal  solution?  Define  and 
illustrate  the  term  "normality  factor." 

7.  Characterize  in  general  the  reactions  upon  which  volumetric  processes 
may  be  based. 

8.  What  is  an  indicator?    Illustrate  your  answer. 

9.  Discuss  the  advantages  of  the  volumetric  system. 

Neutralization  Methods :  The  Standardization  of  Acids  and  Alkalies. 

1.  Define  in  terms  of  the  theory  of  ionization:  (a)  a  neutral  solution; 
(b)  an  acid  solution ;   (c)  an  alkaline  solution. 

2.  Will  the  solution  resulting  from  mixing  equal  volumes  of  one  tenth 
normal  aqueous  solutions  of  the  following  substances  be  acid,  alkaUne,  or 
neutral:  (a)  hydrogen  chloride  and  sodium  hydroxide;  (6)  hydrogen 
chloride  and  ammonium  hydroxide ;  (c)  acetic  acid  and  sodium  hydroxide  ? 

3.  Give  a  full  explanation  of  case  (c)  above,  writing  all  equations  and 
equilibria. 

4.  What  indicator  should  be  used  in  each  case  in  titrations  involving 
the  combinations  indicated  in  question  2  ?  Explain  fully.  (See  the  section 
on  Indicators  for  Use  in  Alkalimetry  and  Acidimetry.) 

5.  OutHne  procedures  for  the  preparation  of  approximately  half -normal 
solutions  of  hydrochloric  acid  and  sodium  hydroxide. 

6.  Describe  a  method  for  obtaining  the  ratio  between  the  solutions 
referred  to  in  the  preceding  question.  If  20.00  cc.  of  an  acid  solution  re- 
quire 21.46  cc.  of  0.4693-normal  alkali^for  neutralization,  what  is  the  nor- 
maHty  factor  of  the  acid? 

7.  Describe  the  sodium  carbonate  method  for  the  standardization  of  a 
solution  of  hydrochloric  acid.  If  0.5383  g.  of  Na2C03  require  20.15  cc.  of 
the  acid  for  neutralization,  what  is  the  normality  factor  of  the  acid? 


178  QUANTITATIVE  CHEMICAL  ANALYSIS 

8.  If  lo.oo  cc.  of  a  solution  of  hydrogen  chloride  yield  0.7421  g.  of  silver 
chloride,  what  is  the  normality  factor  of  the  acid  ? 

9.  Describe  the  standardization  of  sodium  hydroxide  solution  by  means 
of  potassum  bitartrate.  If  1.179  g.  of  the  latter  require  13.35  cc.  of  the 
alkali  for  neutralization,  what  is  the  normahty  factor  of  the  solution? 

10.  If  20.00  cc.  of  an  acid  are  equivalent  to  21.20  cc.  of  an  alkali,  and  if 
40.00  cc.  of  the  acid  are  added  to  0.6000  g.  of  Na2C03  and  the  resulting  solu- 
tion requires  3.00  cc.  of  the  alkaH  for  neutralization,  what  are  the  normality 
factors  of  the  two  solutions? 

11.  Describe  a  method  for  the  preparation  of  a  solution  of  hydrochloric 
acid  of  exactly  one  half  normal  concentration. 

The  Alkaline  Value  of  Soda. 

1.  Discuss  the  composition  of  crude  soda,  and  account  for  the  impurities 
it  is  likely  to  contain. 

2.  Which  of  the  impurities  will  contribute  to  the  alkaline  strength  of 
the  soda? 

3.  Could  this  determination  be  made  with  phenolphthalein  as  the  indi- 
cator, and  if  so  how? 

4.  Give  two  volumetric  methods  for  the  determination  of  Na2C03  and 
of  NaOH  in  mixtures  of  the  two  substances. 

5.  Give  a  method  for  the  determination  of  NazCOa  and  of  NaHCOa  in 
mixtures  of  the  two  salts. 

The  Available  Hydrogen-ion  in  an  Acid. 

1.  In  titrating  weak  acids,  why  is  it  usually  necessary  to  work  at  the 
boiling  temperature? 

2.  At  what  point  does  an  ice-cold,  dilute  solution  of  sodium  hydroxide 
containing  phenolphthalein  lose  its  red  color  upon  being  treated  with  carbon 
dioxide?  Upon  boiling  this  decolorized  solution,  the  color  reappears; 
explain  fully  the  mechanism  by  which  the  alkalinity  of  the  solution  increases 
on  boiling. 

3.  Describe  a  method  for  the  preparation  of  a  carbonate-free  solution 
of  sodium  hydroxide.  How  should  such  a  solution  be  preserved  ?  Why  is 
its  preparation  sometimes  worth  while  ? 

4.  Is  it  possible  to  accurately  titrate  sulphurous  acid  with  a  standard 
solution  of  ammonia?  If  not,  why?  And  if  so,  what  indicator  should  be 
used? 

The  Determination  of  Protein  Nitrogen  by  the  Kjeldahl  Method. 

I.  What  is  the  purpose  in  the  digestion  of  (a)  the  concentrated  sul- 
phuric acid?     (6)  The  copper  sulphate  ?     (c)  The  potassium  sulphate  ? 


QUESTIONS  179 

2.  What  chemical  change  does  sulphuric  acid  undergo  during  the  di- 
gestion?   The  organic  matter? 

3.  Why  is  a  long-necked  flask  used? 

4.  How  would  the  procedure  of  digestion  be  modified  if  nitrates  were 
present?    Explain  fully. 

'  5.  How  would  the  procedure  be  modified  if  mercury  were  added  instead 
of  copper  sulphate  ?    Explain  in  full. 

6.  In  what  form  does  the  nitrogen  exist  after  the  completion  of  the 
digestion?  After  making  the  solution  alkaline  with  sodium  hydroxide? 
In  what  form  does  it  distill  over? 

7.  How  is  the  loss  of  ammonia  prevented  upon  the  addition  of  an  excess 
of  sodium  hydroxide  ? 

8.  What  is  the  purpose  of  the  zinc?  What  is  the  action  of  sodium 
hydroxide  solution  upon  zinc  ? 

9.  The  ammonia  from  one  gram  of  a  fertilizer  is  distilled  into  20.00  cc. 
of  0.5000  N  acid,  and  6.00  cc.  of  0.4800  N  alkali  are  required  to  neutralize 
the  excess  of  acid ;  calculate  the  percentage  of  nitrogen  in  the  sample. 

10.  If  0.20  cc.  of  25.00%  NaOH  (sp.  gr.,  1.25)  had  been  carried  over  by 
bumping  or  foaming,  what  would  have  been  the  apparent  percentage  of 
nitrogen  in  the  above  case  ? 

11.  If  you  had  to  determine  the  percentage  of  NH3  and  of  HC2H3O2  in 
crude  ammonium  acetate,  how  would  you  proceed? 

Dichromate  Methods :  The  Titration  of  Iron. 

1.  Write  the  equation  for  the  oxidation  of  ferrous  chloride  with  potas- 
sium dichromate  in  the  presence  of  hydrochloric  acid. 

2.  What  weight  of  K2Cr207  is  required  for  one  hter  of  a  tenth-normal 
solution,  to  be  used  as  an  oxidizing  agent  ? 

3.  Outline  the  procedure  for  the  standardization  of  dichromate  solu- 
tion by  means  of  iron  wire.  Why  is  it  well  also  to  have  a  standard  solution 
of  ferrous  ammonium  sulphate,  and  how  is  this  solution  standardized  ? 

4.  What  is  the  maximum  weight  of  pure  iron  wire  which  can  be  taken 
for  reaction  with  tenth-normal  dichromate  without  having  to  refill  a  30-cc. 
burette? 

5.  Name  four  reagents  which  can  be  used  to  reduce  ferric  salts  to 
ferrous,  in  the  presence  of  hydrochloric  acid,  and  write  the  equation  in 
each  case. 

6.  Why  is  it  necessary  after  reduction  with  stannous  chloride  to  add 
mercuric  chloride  to  the  solution?  Why  must  the  stannous  chloride  be 
present  only  in  very  slight  excess  (equations)  ?  Why  should  hydrochloric 
acid  be  present  during  the  titration? 


i8o  QUANTITATIVE  CHEMICAL  ANALYSIS, 

7.  What  indicator  is  used  in  the  titration?  What  action  has  it  upon 
ferric  salts?    Upon  ferrous  salts? 

8.  Why  is  the  indicator  used  outside  of  the  solution? 

9.  Can  you  give  a  method  of  treatment  for  chrome  iron  ore  which  might 
lead  to  a  determination  of  its  chromium  content  by  means  of  standard 
solutions  of  ferrous  iron  and  dichromate? 

Permanganate  Methods :  The  Titration  of  Iron  and  of  Oxalic  Acid. 

1.  How  many  grams  of  KMn04  are  required  for  one  liter  of  tenth-normal 
solution,  to  be  used  as  an  oxidizing  agent  in  acid  solutions  ?  To  be  used  in 
neutral  solutions  for  the  oxidation  of  manganese  ? 

2.  Why  should  a  permanganate  solution  be  allowed  to  stand  for  several 
days,  and  then  be  filtered  through  asbestos,  before  it  is  standardized? 
Why  should  it  not  be  placed  in  burettes  having  rubber  outlet  tubes  ? 

3.  Name  at  least  four  substances  which  can  be  used  to  standardize 
permanganate  solutions.    Write  an  equation  in  each  case. 

4.  What  is  the  maximum  weight  of  Na2C204  which  can  be  titrated  with 
tenth-normal  permanganate  solution  without  having  to  refill  a  30-cc.  burette? 

5.  In  the  titration  of  oxalic  acid,  why  is  it  that  the  oxidation  proceeds 
so  much  more  slowly  at  first  than  later  on?    Explain  in  full. 

6.  Name  ten  substances  which  can  be  quantitatively  determined  by 
means  of  potassium  permanganate. 

7.  What  effect  does  potassium  permanganate  have  upon  hydrochloric 
acid  in  the  presence  of  ferrous  salts,  even  in  very  dilute  solution?  How 
may  this  action  be  prevented  ? 

8.  What  are  the  components  of  the  Zimmermann-Reinhardt  solution? 
Explain  the  purpose  of  each. 

9.  Under  what  conditions  can  ferrous  iron  be  accurately  determined 
with  potassium  permanganate  without  the  use  of  the  Zimmermann-Rein- 
hardt solution  ? 

10.  What  is  the  maximum  weight  of  a  sample  of  ore  containing  40.00% 
of  iron  which  can  be  taken  for  titration  with  tenth-normal  oxidizing  agent 
without  having  to  refill  a  30-cc.  burette  ? 

11.  Discuss  the  preparation  of  a  solution  for  analysis  from  a  refractory 
iron  ore. 

12.  Describe  a  method  for  the  determination  of  calcium  by  means  of 
potassium  permanganate. 

13.  Outline  a  method  for  the  determination  of  the  MnOrvalue  of 
pyrolusite  by  means  of  potassium  permanganate. 

14.  How  can  the  determination  referred  to  in  the  preceding  question  be 
made  by  a  gravimetric  process? 


QUESTIONS  i8i 

The  Determination  of  Phosphorus  in  Steel. 

1.  Assuming  the  presence  of  the  phosphorus  as  Fe3P2,  show  by  means 
of  an  equation  the  action  of  nitric  acid  upon  this  compound.  Why  is  the 
nitric  acid  solution  heated  with  potassium  permanganate? 

2.  In  order  to  cause  the  higher  oxides  of  manganese,  such  as  Mn02,  to 
go  into  solution  in  nitric  acid,  what  kind  of  a  reagent  should  be  added? 
Illustrate  and  explain. 

3.  What  is  the  purpose  of  precipitating  the  phosphorus  as  anmionium 
phosphomolybdate  ? 

4.  Why  is  the  yellow  precipitate  dissolved  in  ammonia  and  the  solution 
acidified  with  sulphuric  acid,  rather  than  to  dissolve  it  directly  in  sulphuric 
acid?  Why  is  M0O3  not  precipitated  when  the  ammoniacal  solution  is 
acidified  with  sulphuric  acid  ?  Would  hydrochloric  or  nitric  acid  do  as  well 
here,  and  why? 

5.  Describe  the  construction  and  use  of  the  Jones  reductor.  Can  it  be 
used  for  the  reduction  of  substances  other  than  molybdenum  ? 

6.  Why  is  it  best  to  receive  the  reduced  molybdenum  solution  below 
the  surface  of  a  solution  containing  ferric  alum?  What  is  the  purpose  of 
the  phosphoric  acid  in  this  solution  ? 

7.  How  should  the  permanganate  solution  be  standardized  in  order  to 
obtain  the  most  reliable  results  ? 

8.  Would  you  recommend  the  determination  of  phosphoric  anhydride 
in  apatite  by  this  method  ?  Why?  Is  there  a  suitable  volumetric  method ? 
If  so,  describe  it. 

The  Determination  of  Manganese  in  an  Ore. 

1.  What  is  the  fundamental  reaction  of  this  process?  What  r61e 
does  it  play  in  the  titration  of  oxalic  acid  or  of  iron  with  potassium 
permanganate  ? 

2.  Discuss  the  preparation  of  the  solution  for  analysis  from  a  refractory 
ore  containing  manganese. 

3.  What  happens  when  zinc  oxide  is  added  to  the  acid  solution  of  the 
ore  (equations)  ? 

4.  Explain  fully  why  a  zinc  salt  should  be  present  in  the  solution  during 
the  titration. 

5.  Explain  why  the  presence  of  chlorides  does  not  interfere  with  the 
accuracy  of  this  titration. 

6.  If  the  permanganate  solution  used  in  this  titration  is  o.iooo-iV  for 
use  with  iron  or  oxalic  acid,  what  is  its  normaUty  factor  for  this  reaction? 

7.  How  is  it  best  to  standardize  the  permanganate  solution  used  in  this 
determination? 


i82  QUANTITATIVE  CHEMICAL  ANALYSIS 

lodometric  Methods :  The  Preparation  and  Standardization  of  Iodine 
and  Thiosulphate  Solutions. 

1.  How  many  grams  of  iodine  are  required  for  one  liter  of  the  tenth- 
normal solution?    Of  sodium  thiosulphate,  Na2S203 .  5  H2O? 

2.  How  is  the  iodine  solution  made,  and  what  is  the  purpose  of  the 
potassium  iodide?  Show  what  equilibria  exist  in  the  iodine  solution.  Why 
should  the  water  be  freshly  boiled  and  allowed  to  cool  out  of  contact  with 
the  air,  in  the  preparation  of  the  thiosulphate  solution? 

3.  Write  the  reaction  between  iodine  and  sodium  thiosulphate.  How 
do  chlorine  and  bromine  differ  from  iodine  in  their  behavior  towards  sodium 
thiosulphate  ?    Explain  why  this  is  so. 

4.  What  is  the  effect  of  free  carbonic  acid  upon  sodium  thiosulphate 
solution?  Does  the  decomposition  cease  as  soon  as  all  of  the  carbonic 
acid  has  reacted  ?    Write  equations  to  illustrate  your  answers. 

5.  Does  the  solution  resulting  from  the  partial  decomposition  of  the 
thiosulphate  have  a  greater  or  lower  reducing  value  than  the  original  solu- 
tion?   Explain  why. 

6.  Give  equations  to  show  two  ways  in  which  iodine  may  act  as  an 
oxidizing  agent. 

7.  What  is  the  maximum  weight  of  AS2O3  which  can  be  taken  for 
reaction  with  tenth-normal  iodine  solution  without  having  to  refill  a  30-cc. 
burette  ? 

8.  Can  the  standardization  of  iodine  against  arsenious  oxide  be  per- 
formed in  a  strongly  alkaline  solution  ?  Can  it  be  done  in  an  acid  solution  ? 
Give  reasons  for  your  answers. 

9.  What  is  the  purpose  of  the  sodium  bicarbonate  ?  Is  the  bicarbonate 
solution  acid,  alkaline,  or  neutral  ?    Explain  your  answer. 

10.  Discuss  the  determination  of  the  end-point.  Explain  why  the  indi- 
cator is  added  in  such  large  quantity. 

11.  Discuss  the  use  of  iodine  solutions  in  the  presence  of  sulphuric  acid. 
In  the  presence  of  ammonium  salts. 

The  Determination  of  Antimony  in  Stibnite. 

1.  Write  the  reaction  between  pure  stibnite  and  hydrochloric  acid. 

2.  Why  must  the  hydrochloric  acid  solution  be  heated  on  the  steam 
bath  ?    Why  must  it  not  be  boiled  until  after  dilution  ? 

3.  What  is  the  purpose  of  adding  tartaric  acid  to  the  solution  ?    Explain. 

4.  Explain  why  the  solution  may  possibly  turn  red  during  gradual 
dilution.     What  is  the  correct  procedure  in  such  a  case  ? 

5.  If  a  white  precipitate  forms  upon  dilution,  what  error  has  been  made? 
What  is  the  white  precipitate,  and  what  should  be  done  with  the  mixture  ? 


QUESTIONS  183 

6.  What  is  the  purpose  of  almost  neutralizing  the  solution  with  sodium 
hydroxide,  and  how  is  this  accomphshed?  Why  is  sodium  bicarbonate 
then  added  in  large  excess? 

7.  What  elements  would,  if  present,  interfere  with  this  determination, 
and  why? 

The  Determination  of  Chromium  in  Chromite. 

1.  What  is  the  action  of  fused  sodium  peroxide  upon  ferrous  oxide? 
Upon  chromic  oxide  ? 

2.  What  happens  upon  the  addition  of  water  to  the  cooled  chromite- 
sodium  peroxide  fusion  mixture  ? 

3.  Why  is  it  necessary  to  completely  decompose  the  excess  of  the  peroxide 
before  acidifying  the  aqueous  extract  ? 

4.  Give  two  processes  for  the  volumetric  determination  of  chromium  in 
the  acidified  extract,  and  explain  each. 

The  Determination  of  Lead  in  an  Ore. 

1.  After  the  decomposition  of  the  ore  and  the  addition  of  sulphuric  acid, 
why  is  it  necessary  to  evaporate  to  white  fumes  ? 

2.  Explain  the  solubility  of  lead  sulphate  in  ammonium  acetate  solution. 

3.  Explain  why  it  is  possible  to  quantitatively  precipitate  the  lead  from 
the  ammonium  acetate  solution  by  means  of  an  excess  of  potassium  dichro- 
mate. 

4.  Explain  the  solubihty  of  lead  chromate  in  the  acidified  solution  of 
sodium  chloride. 

5.  Write  an  equation  to  show  the  reaction  of  the  acid  chromate  solution 
with  potassium  iodide. 

6.  Why  should  the  thiosulphate  solution  used  in  this  determination  be 
standardized  under  identical  conditions  against  test  lead  ? 

The  Determination  of  Copper  in  an  Ore. 

1.  In  this  determination,  why  is  the  copper  separated  from  the  other 
metals  present  in  the  ore?  Explain  how  iron,  arsenic,  or  antimony  would 
interfere  with  the  accuracy  of  the  titration. 

2.  What  would  you  expect  the  composition  of  the  precipitate  to  be 
which  is  formed  upon  the  addition  of  sodium  thiosulphate  to  a  solution  of 
copper  sulphate?  How  does  it  happen,  then,  that  we  obtain  cuprous 
sulphide  ? 

3.  Can  any  other  metals  be  precipitated  from  their  salt  solutions  by 
means  of  sodium  thiosulphate?  (Try,  for  example,  silver  nitrate  and 
sodium  thiosulphate,  in  the  cold,  and  explain  what  takes  place.) 


1 84  QUANTITATIVE  CHEMICAL  ANALYSIS 

4.  What  is  the  object  of  igniting  the  precipitate  of  cuprous  sulphide? 
What  becomes  of  any  antimony  which  is  present  ? 

5.  Why  is  it  so  important  to  standardize  the  thiosulphate  solution 
against  pure  metallic  copper? 

6.  Why  must  the  nitrous  fumes  be  completely  expelled  before  the  ad- 
dition of  the  potassium  iodide  ? 

7.  Why  is  it  preferable  to  titrate  the  free  iodine  in  the  presence  of  acetic 
acid,  rather  than  in  the  presence  of  sulphuric  acid  ?    Explain  fully. 

8.  Write  an  equation  to  show  the  action  of  nitric  acid  upon  cuprous 
iodide. 

Precipitation  Methods :  The  Determination  of  Chlorine. 

1.  Briefly  outline  the  procedure  for  the  standardization  of  the  silver 
nitrate  and  potassium  thiocyanate  solutions  against  pure  sodium  chloride. 
Could  these  solutions  be  standardized  against  pure  metallic  silver,  and  if 
so  how? 

2.  How  may  pure  sodium  chloride  be  prepared  from  the  commercial  salt  ? 

3.  What  indicator  is  used  in  connection  with  thiocyanate  solutions? 
Why  must  nitric  acid  be  present  ?  Explain  fully  why  the  indicator  should 
be  added  in  such  large  quantity. 

4.  At  18-20°,  the  solubility  product  of  silver  chloride  is  about  o.6Xio~io 
and  that  of  silver  thiocyanate  is  about  o.6Xio~^'^;  what  is  the  relative  con- 
centration of  the  chloride  and  thiocyanate  ions  in  a  solution  which  is  satu- 
rated with  both  salts?  Assume  the  equal  (practically  complete)  ioniza- 
tion of  both  salts. 

5.  Why  is  it  necessary  to  filter  off  the  silver  chloride  before  making  the 
titration  with  the  thiocyanate  solution?  Base  your  explanation  upon  the 
data  given  in  the  preceding  question. 

6.  In  general,  what  anions  may  be  determined  by  this  method  without 
first  filtering  off  the  silver  salt  ? 

7.  How  may  the  halogens  in  alkali  chlorates,  bromates,  and  iodates  be 
determined  by  this  metjiod  ? 

8.  Outline  a  procedure  for  the  determination  by  this  method  of  {a)  the 
chlorine  in  horn  silver,  AgCl ;  (Jb)  the  silver  in  the  same  mineral. 


PART   V 

ANALYTICAL  PROBLEMS 

Preliminary  Discussion.  In  spite  of  any  preconception  to 
the  contrary,  the  calculations  involved  in  analytical  work  are 
not  difficult,  and  a  reasonable  amount  of  application  will  enable 
the  student  to  master  the  subject.  At  the  same  time,  the  study 
of  the  following  problems  will  furnish  an  insight  into  the  prin- 
ciples of  a  wide  variety  of  processes. 

It  cannot  he  too  strongly  emphasized  that,  in  making  analyti- 
cal calculations,  the  beginner  should  from  the  outset  strive  to  take 
the  shortest  and  most  direct  route  to  the  result.  With  a  little 
practice,  the  student  who  is  not  unacquainted  with  the  reactions 
of  analytical  chemistry  should  soon  be  able  upon  the  inspection 
of  a  problem  to  promptly  recognize  the  factors  which  will  lead 
most  directly  to  its  solution,^  as  well  as  the  equivalent  relation- 
ships of  the  substances  involved.     To  do  this,  however,  he  should 

*  Of  course  most  analytical  problems  can  be  solved  in  stages,  by  means  of  a 
series  of  proportions,  and  it  is  perhaps  only  natural  that  most  beginners  should 
have  a  predilection  for  this  method.  In  the  examples  given,  however,  the  common 
factors  have  been  eliminated,  and  the  problems  solved  in  a  single  operation.  The 
beginning  student  will  better  appreciate  the  advantages  of  the  shorter  method 
upon  comparing  the  solutions  given  of  problems  iv  and  v  with  the  following 
roundabout  method  of  arriving  at  the  same  results : 

iv.  (a)  lo  Fe :  2  KMn04= 0.005  :  w. 

^=316/558X0.005 =0.00283  g-  KMn04  per  cc. 
(6)  2  KMn04 :  5  H2C2O4 = 0.00283 '  x. 

a;=45o/3i6Xo.oo283  =  o.oo403  g.  H2C2O4  per  cc. 

(c)  5  H2C2O4 :  5  CaC204 = 0.00403 :  y. 
3^=640/450X0.00403=0.00573  g.  CaC204  per  cc. 

(d)  5  CaC204 :  5  CaO= 0.00573  :  z. 
2=280.5/640.5X0.00573=0.00251  g.  CaO  per  cc.    Arts. 


i86  QUANTITATIVE  CHEMICAL  ANALYSIS 

understand  and  hear  in  mind  the  relationships  and  differences 
which  exist  between  chemical  and  physical  units,  —  such  as  atoms , 
molecules,  and  equivalents  on  the  one  hand,  and  grams  and  cubic 
centimeters  on  the  other. 

In  order  to  illustrate  what  is  meant,  the  following  problems 
are  given,  with  their  solutions. 

i.  A  sample  of  a  soluble  chloride  weighing  0.2007  g.  yields  on 
analysis  0.4920  g.  of  silver  chloride;  what  percentage  of  chlorine 
does  it  contain? 

From  the  proportion, 

Wt.  of  chlorine  in  sample :  Wt.  of  sample  =  %  of  chlorine :  100, 

,    .        ,,    ,  Wt.  of  chlorine  in  sample  ^     r    1.1    • 

it  IS  obvious  that — ; — ^—  •  100  =  %  of  chlorine. 

Wt.  of  sample 

Also,  since  the  chlorine  contained  in  the  sample  is  identical  with 

that  which  is  later  contained  in  the  silver  chloride  precipitate, 

we  have  the  proportion, 

Cl:AgCl=Wt.  of  chlorine  :Wt.  of  silver  chloride, 

CI 

or,  •  Wt.  of  silver  chloride =Wt.  of  chlorine. 

AgCl 

Substituting  this  value  in  the  preceding  equation,  we  get, 

-^—'  Wt.  of  silver  chloride  35-46  .  0.4920 

AgCl 143-34 

Wt.  of  sample  '  ^^~        0.2007       *  ^°° 

=  60.50%  of  CI. 

A  chemical  factor  represents  the  quantity  by  weight  of  an  ele- 
ment or  compound  which  is  equivalent  to  one  part  by  weight  of 
some  other  substance.     For  example,  the  ratio  or  factor 

Ag       107.88  ^ 

.  ° ,  =  — - —  =  0.7526 
AgCl     143.34       '^ 

V,  (o)  AgCl:HCl=o.i527::c. 

a;  =  36.46/143.34X0.1527  =  0.0388  g.  HCl  in  20.50  cc. 
{h)  20.50 :  1000=0.0388  :  y. 

3;  =  1000/20.5X0.0388=  1.893  g.  HCl  in  one  liter, 
(c)  36.46:1.893  =  1:2. 

2=1.893/36.46=0.0519  iV.    Ans. 


ANALYTICAL  PROBLEMS  187 

tells  us  that  one  gram  of  silver  chloride  contains  0.7526  g.  of 
silver,  and  if  we  wish  to  calculate  what  weight  of  silver  there  is 
in  a  specific  weight  of  silver  chloride,  we  simply  multiply  the 
latter  by  this  factor;  e.g.  10.15  g.  of  silver  chloride  contain 
10.15X0.7526=7.64  g.  of  silver. 

Or,  if  the  weight  of  MnO  is  wanted  which  corresponds  to  a 
definite  weight  of  Mn304,  the  factor  is 

S  MnO      212.70 

— = ^  =  0.0'^0I. 

MnaOi      228.79 

And  if  we  wish  to  find  the  weight  of  K2O  which  corresponds  to  a 
given  weight  of  KCIO4,  we  multiply  the  latter  by  the  factor, 

K2O         04.20 

Since,  in  arriving  at  these  (physical-unit)  factors,  our  calcu- 
lations must  be  based  upon  chemical  facts,  it  is  necessary  to  keep 
strictly  to  the  equivalent  relations  of  the  substances.  Thus,  in 
the  last  example,  since  every  molecule  of  K2O  would  be  converted 
upon  evaporation  with  perchloric  acid  into  two  molecules  of 
KCIO4,  we  must  use  one  mol  of  K2O  in  the  numerator  and  two 
of  KCIO4  in  the  denominator.  Any  other  ratio  would  Jie^in- 
correct. 

ii.  How  many  cubic  centimeters  of  a  solution  containing  12.5 
grams  of  BaCh  -  2  H2O  per  liter  will  he  required  to  precipitate 
the  sulphur  from  0.1073  gram  of  pure  stibnite,  Sb^Sz,  as  BaSOi? 

Each  molecule  of  Sb2S3  will  yield  upon  treatment  three  mole- 
cules of  H2SO4,  and  these  will  react  with  three  molecules  of  barium 
chloride.  We  therefore  arrive  at  the  proportion,  3(BaCl2- 
2H2O) :  Sb2S3  =  x :  0.1073,  where  x  is  the  weight  of  the  crystalline 
salt  required.     That  is, 

3(BaCl2.  2  H2O) 
^"^    cue  '  o-i°73, 

in  which  the  factor  ^^   ^cuo^ — ^  =  2.177  indicates  the  quan- 

00203 


■\ 


i88  QUANTITATIVE  CHEMICAL  ANALYSIS 

tity  by  weight  of  BaCl2 .  2  H2O  which  is  required  to  precipitate 
the  sulphur  from  one  gram  of  Sb2S3;  this  ratio  therefore 
does  not  differ  essentially  from, the  chemical  factors  previously 
discussed. 

Finally,  since  each  cubic  centimeter  of  the  solution  contains 
0.0125  g.  of  BaCl2 .  2  H2O,  we  have, 

3(BaCl2 .  2  H2O)  ^         732.9 

Sb2S3 ^  336:6 ^  ^g  ^g  ^^^ 

0.0125  0.0125 

iii.  What  volume  of  aqueous  ammonia  of  sp.  gr.  0.960,  contain- 
ing 9.91%  of  NHz,  will  be  required  to  precipitate,  as  Fe{OB)z,  the 
iron  contained  in  1.475  g.  of  Fe{NH^S0^2  -  6  H2O? 

Since  the  iron  is  to  be  precipitated,  after  oxidation,  as  Fe(0H)3, 
we  see  that  each  atom  of  iron  will  require  three  molecules  of 
NH4OH,  which  in  turn  are  furnished  by  three  molecules  of  NH3. 

and,  since  each  cubic  centimeter  of  the  aqueous  ammonia  weighs 
0.960  g.,  and  contains  9.91%  of  NH3,  the  solution  contains 
0.960X0.0991  g.  of  NH3  per  cubic  centimeter.    That  is, 

Fe(NH4S04)2.6H,0      ^^'_392-i6  _;.o;  cc. 

0.960X0.0991  0.960X0.0991 

iv.  A  solution  of  potassium  permanganate  is  equivalent  to 
0.00500  g.  of  ferrous  iron  per  cubic  centimeter;  what  is  its  value 
in  terms  of  calcium  oxide  ? 

The  reactions  involved  in  the  volumetric  determinations  of 
iron  and  calcium  are : 

10  FeS04+2  KMn04+9  H2S04=  2  KHSO4+2  MnS04 

+5  Fe2(S04)3+8  H2O, 
and 

5  CaC204+2  KMn04+9  H2S04=5  CaS04+2  KHSO4 

+  2MnS04+ioC02+8H20; 


ANALYTICAL  PROBLEMS  189 

and  from  these  equations  it  is  seen  that,  in  this  case, 
2Fe++=c=CaO.    That  is, 


CaO 


•  0.00500=   ^      •0.00500=0.00251  g.  CaO. 


2  Fe  111.6 

V,  If  20.50  cc.  of  hydrochloric  acid  yield  0.1527  g.  of  silver 
chloride  J  what  is  the  normality  factor  of  the  solution? 

Although  silver  chloride  is  insoluble,  the  normahty  factor  of 
the  acid  may  nevertheless  be  calculated  directly  from  the  weight 
of  the  precipitate  obtained.  One  liter  of  normal  hydrochloric 
acid,  containing  one  mol  of  HCl,  would  yield  143.34  g.  {i.e.  one 
mol)  of  silver  chloride,  whence  20.50  cc.  would  yield  0.14334 
X  20.50  g.    We  obtain,  therefore,  the  equation, 

^lii^T =  0.0519  iV^. 

0.14334X20.50 

vi.  A  sample  of  stihnite  weighing  0.1793  g.  is  heated  with  strong 
ECl,  and  the  H2S  evolved  absorbed  by  means  of  sodium  hydroxide 
solution;  the  resulting  mixture  (containing  the  sulphur  as  Na2S) 
being  introduced  under  the  surface  of  a  solution  made  by  adding 
25  cc.  of  6-normal  HCl  and  50.00  cc.  of  0.1160-normal  iodine 
to  500  cc.  of  water.  The  excess  of  iodine  is  titrated  with  0.0957- 
normal  sodium  thiosulphate  solution ^  of  which  28.57  ^^-  ^^^  ^^- 
quired.    Calculate  the  percentage  of  (evolved)  sulphur  in  the  stibnite. 

(H2S  +  l2=2HI-l-S.) 

cc.        N.  F. 

50.00X0.1160=5.800  cc.  of  normal  iodine. 
28.57X0.0957=  2.734  cc.  of  normal  thiosulphate. 

I.e.  the  H2S  required     3.066  cc.  of  normal  iodine. 

Since  normal  iodine  has  a  sulphur  value  of  ^^^^^52= 0.016035  g. 

2 

per  cubic  centimeter,  we  have, 

o.oi6os';X^.o66^^  /w    r     11 
^^ — ■^ X 100=  27.42%  of  sulphur. 

0.1793 


IQO  QUANTITATIVE  CHEMICAL  ANALYSIS 

The  normality  factor  of  a  solution  expresses  the  value  of  the 
solution  per  cubic  centimeter  in  terms  of  a  normal  solution. 
For  example,  if  a  solution  is  known  to  be  one  half  normal  {i.e. 
N.  F.  =  0.500),  it  is  obvious  that  i  cc.  of  it  is  equivalent  to 
1.000X0.500=0.500  cc.  of  a  normal  solution;  or  that  27.31  cc. 
of  it  are  equivalent  to  27.31X0.500=13.655  cc.  of  a  normal 
solution.  Knowing  the  normality  factors  of  a  series  of  solutions, 
therefore,  we  can  readily  reduce  the  different  volumes  of  the 
solutions  used  in  a  determination  to  a  common  standard,  and 
in  this  way  render  the  calculations  quite  as  simple  as  if  the  solu- 
tions had  all  been  made  up  to,  say,  exactly  one  tenth  normal. 

vii.  A  mixture  of  potassium  chloride  and  caesium  chloride 
weighs  0.3895  g.  This  is  dissolved  in  water  and  the  chlorine  pre- 
cipitated as  silver  chloride,  of  which  there  is  obtained  0.4889  g. 
Calculate  the  weight  of  each  salt  in  the  mixture. 

Indirect  methods  of  analysis  depend  upon  the  fact  that  when 
two  or  more  substances  are  made  to  undergo  the  same  chemical 
treatment  they  either  experience  a  relatively  different  change  of 
weight,  or  unit  weights  of  each  require  unequal  volumes  of  a 
standard  solution. 

Let  X  represent  the  weight  of  the  potassium  chloride,  in  the 
above  example,  and  y  that  of  the  caesium  chloride,  a  the  weight 
of  the  two  salts,  and  p  that  of  the  precipitate,  and  we  have, 

x-[-y=a, 
and  AgCl       AgQ 

KCl        CsCl  ^    ^ 

If  we   designate  by  m  the  factor  ^      and  by  n  the  factor 

KCl 

— %— ,  we  obtain, 

CsCr  '  x-^y=a, 

and  mx-\-ny=py 

from  which  we  find  that 

p—na 

x=- , 

m—n 

1        .        n 

or  x= •^  — 


m—n         m—n 


ANALYTICAL  PROBLEMS  191 

Indirect  analyses  may  in  general  be  calculated  by  means  of 
this  or  a  similar  general  equation. 
In  the  above  example, 

KCl      74.56  CsCl     168.27 

and  w— w=i.o7io. 

Substituting  these  values  in  the  general  equation,  we  obtain, 
x=-o.gss7  p -0.7949  a. 

Consequently,  in  order  to  determine  the  weight  of  potassium 
chloride  in  the  mixed  sample  it  is  only  necessary  to  multiply 
the  values  of  a  and  p  by  0.7949  and  0.9337,  respectively,  and 
subtract  the  first  product  from  the  second ;  i.e., 

ix^=o.9337Xo.4889-o.3895Xo.7949  =  o.i47o  g.  KCl, 

and    ^=0.3895 -0.1470=0.2425  g.  CsCl. 

The  same  analysis  might  be  performed  by  weighing  the  mixed 

chlorides  in  a  platinum  crucible,  then  changing  them  to  sulphates 

(by  treatment  with  H2SO4,  etc.),  and  again  weighing.    In  this 

case  also,  ^  ^ 

x=-l—'P--^^'a 
m—n         m—n 

=  10.66  /»  — 11.45  a. 

In  the  first  case,  the  coefficients  are  relatively  small,  and  con- 
sequently good  results  might  be  expected,  since  the  experimental 
errors  made  in  the  determination  of  a  and  p  are  multiplied  by 
only  0.7949  and  0.9337,  respectively.  In  the  latter  case,  how- 
ever, the  coefficients  are  very  large,  and  the  unavoidable  analyti- 
cal errors  would  have  to  be  multiplied  enormously  in  the  cal- 
culation ;  the  latter  method  is  therefore  worthless. 

Although  some  indirect  methods  may  appear  simple  and 
attractive  on  paper,  they  frequently  lead  to  impossible  values  in 
practice ;  so  that  extreme  caution  should  be  exercised  regarding 
the  use  of  an  indirect  method.  In  general,  if  accurate  and  reliable 
results  are  desired,  indirect  methods  of  analysis  should  he  avoided. 


192  QUANTITATIVE  CHEMICAL  ANALYSIS 

PROBLEMS 
GRAVIMETRIC  ANALYSIS 

1.  If  the  determination  of  a  substance  by  a  certain  method  permits  of 
an  error  of  o.i%  in  the  weighing  of  the  sample,  how  accurately  must  a  200 
mg.  sample  be  weighed?    A  5  g.  sample? 

2.  Calculate  the  chemical  factors  for  the  following :  KCl  from  KaPtCle ; 
K20fromKC104;  MgO  f rom  Mg2P207 ;  Mn  f rom  Mn2P207 ;  MnOa  from 
Mn304. 

3.  What  weight  of  Mn304  is  equivalent  to  0.5785  g.  of  Mn2P207?  To 
04327  g.  of  MnS04? 

4.  A  portion  of  phosphorus  pent  oxide  weighing  0.2018  g.  yields  0.3132  g. 
of  Mg2P207.     Calculate  the  percentage  of  P2O5  in  the  sample. 

5.  A  0.4988  g.  sample  of  a  salt,  upon  distillation  with  sodium  hydroxide 
solution  evolves  ammonia,  which  is  converted  into  (NH4)2PtCl6;  and  this 
upon  ignition  yields  0.3258  g.  of  platinum.  Calculate  directly  from  the 
weight  of  platinum  the  percentage  of  NH3  in  the  sample. 

6.  A  quantity  of  the  silver  salt  of  an  organic  acid  weighing  0.4072  g. 
yields  upon  ignition  0.2632  g.  of  metallic  silver.  Calculate  the  percentage 
of  silver  in  the  salt. 

7.  What  weight  of  a  silver  nitrate  solution  known  to  contain  2.31% 
of  Ag  will  be  required  to  precipitate  the  chlorine  from  25.0  cc.  of  absolution 
containing  12.5  g.  of  BaCl2 .  2  H2O  in  one  Hter? 

J8.  If  25.0  cc.  of  sodium  chloride  solution  yield  0.1434  g.  of  silver  chloride, 
what  is  the  strength  of  the  solution  in  grams  of  the  salt  per  liter?  In  mols 
per  liter? 

g.  How  many  cubic  centimeters  of  hydrochloric  acid  of  sp.  gr.  1.050, 
containing  10.17%  of  HCl,  will  it  take  to  precipitate  the  silver  from  a  solu- 
tion containing  0.8430  g.  of  silver  sulphate  ? 

10.  How  many  cubic  centimeters  of  hydrochloric  acid  of  sp.  gr.  1.040, 
containing  8.16%  of  HCl,  will  be  required  to  dissolve  one  gram  of  calcium 
carbonate  ? 

11.  What  weight  of  Mn2P207  is  it  possible  to  prepare  from  50.0  cc.  of  a 
permanganate  solution  which  contains  4.500  g.  of  KMn04  per  liter? 

12.  How  many  cubic  centimeters  of  a  solution  of  sp.  gr.  1.116,  con- 
taining 10.06%  of  NaOH,  will  it  take  to  neutralize  a  solution  containing 
5.00  g.  of  NaHS04?  5.00  g.  of  KHSO4? 

13.  A  sample  of  impure  potassium  sulphide  weighing  0.4320  g.  is  treated 
with  hydrochloric  acid,  and  by  means  of  ammoniacal  hydrogen  peroxide 
solution  the  hydrogen  sulphide  evolved  is  converted  into  ammonium  sul- 


ANALYTICAL  PROBLEMS  193 

phate.    This  yields  0.8034  g.  of  BaS04.    Calculate  the  percentage  of  K2S 
in  the  sample. 

14.  A  sample  of  stibnite  weighing  1.078  g.,  upon  being  analyzed  by  the 
method  indicated  in  Problem  13,  yields  0.6750  g.  of  BaS04.  Assuming 
the  sulphur  to  be  present  wholly  as  Sb2S3,  calculate  the  percentage  of  the 
latter  in  the  mineral. 

15.  How  many  cubic  centimeters  of  aqueous  ammonia  of  sp.  gr.  0.96, 
containing  9.91%  of  NH3,  will  be  required  to  precipitate  the  alumimma  in 
0.8674  g.  of  KA1(S04)2 .  12  H2O  ?  How  many  cubic  centimeters  of  6-normal 
ammonia  ? 

16.  What  volume  of  the  ammonia  water  first  referred  to  in  Problem  15 
will  it  take  to  neutralize  lo.o  cc.  of  hydrochloric  acid  of  sp.  gr.  1.12,  con- 
taining 23.81%  of  HCl?  To  neutralize  lo.o  cc.  of  6-normal  hydrochloric 
acid? 

17.  If  15.0  cc.  of  a  solution  of  barium  chloride  5deld,  upon  evaporation 
with  hydrochloric  acid  and  gentle  ignition,  1.563  g.  of  the  anhydrous  salt, 
what  is  the  strength  of  the  solution  in  mols  per  liter?  In  equivalents  per 
liter? 

18.  A  solution  contains  2.25%  by  weight  of  BaCl2.  What  volume 
of  a  solution  containing  5.000  g.  of  Ag2S04  per  liter  will  be  required  to  pre- 
cipitate the  chlorine  in  5.000  g.  of  the  barium  chloride  solution?  What 
will  be  the  weight  of  the  precipitate  formed  ? 

19.  A^sample  of  pyritic  mineral  weighing  0.2637  g-  yields  upon  analysis 
0.7993  g-  of  BaS04.  Upon  the  assumption  that  the  sulphur  is  wholly  pres- 
ent as  FeS2,  calculate  the  percentage  of  the  latter  in  the  mineral. 

20.  What  volume  of  bromine  water  containing  3.0%  by  weight  of  bro- 
mine will  be  required  to  oxidize  the  iron  in  1.75  g.  of  FeS04 .  7  H2O? 

21.  What  volume  of  aqueous  ammonia  (sp.  gr.  0.96,  containing  9.91% 
of  NH3)  will  be  required,  after  oxidation  of  the  iron  with  hydrogen  peroxide, 
to  precipitate  the  metal  from  a  solution  made  by  dissolving  0.750  g.  of 
Fe(NH4S04)2 .  6  H2O  in  water  acidified  with  12.0  cc.  of  hydrochloric  acid 
(sp.  gr.  1. 12,  containing  23.8%  of  HCl)? 

22.  A  mixed  sample  of  CaO,  Ca(0H)2,  and  CaCOs  weighing  0.5896  g. 
is  evaporated  with  excess  sulphuric  acid,  and  gently  ignited ;  the  residue 
weighs  0.8651  g.  What  volume  of  6-normal  hydrochloric  acid  will  be  re- 
quired to  convert  5.00  g.  of  the  sample  into  calcium  chloride? 

23.  A  sample  of  pyrolusite  weighing  0.5124  g.  is  heated  in  the  presence 
of  dilute  sulphuric  acid  with  an  excess  of  oxalic  acid,  and  the  gas  evolved 
is  absorbed  in  a  weighed  bulb  containing  potassium  hydroxide.  The  gain 
in  weight  of  the  bulb  is  found  to  be  0.4789  g.  Calculate  the  percentage  of 
Mn02  in  the  pyrolusite. 


194  QUANTITATIVE  CHEMICAL  ANALYSIS 

24.  The  ignited  mixture  of  ferric  and  aluminum  oxides  from  1.497  g- 
of  a  mineral  weighs  0.4196  g. ;  after  a  second  ignition,  in  a  current  of 
hydrogen,  the  product  weighs  0.331 1  g.,  the  ferric  oxide  being  reduced 
to  metallic  iron.  Calculate  the  percentage  of  Fe203  and  of  AI2O3  in  the 
mineral. 

25.  What  volume  of  0.5-normal  ammonium  oxalate  solution  will  be 
required  to  precipitate,  as  CaC204 .  H2Q,  the  calcium  from  one  gram  of 
apatite,  [Ca3(P04)2]3  •  CaFa? 

26.  What  volume  of  a  solution  containing  66  g.  of  (NH4)2HP04  per  liter 
will  be  required  to  precipitate,  as  ZnNH4P04,  the  zinc  from  0.9786  g.  of  a 
brass  which  contains  30.15%  of  zinc?  What  is  the  normality  of  this  solu- 
tion as  a  precipitant  for  zinc  ? 

27.  How  many  grams  per  liter  of  (NH4)2Cr207  must  a  solution  contain 
in  order  that,  by  reduction  of  the  chromium  (with  HCl  and  SO2),  precipita- 
tion with  ammonia,  and  ignition  of  the  precipitate  in  a  current  of  hydrogen, 
a  25.0  cc.  portion  shall  yield  0.1267  g.  of  Cr203? 

28.  What  volume  of  6-normal  sulphuric  acid  will  be  required  to  replace 
the  nitric  acid  in  the  salts  obtained  upon  evaporating  to  dryness  the  solu- 
tion in  nitric  acid  of  a  five  cent  coin  weighing  4.960  g.,  the  composition 
of  the  coin  being  75.00%  Cu  and  25.00%  Ni? 

29.  A  sample  of  silicate  mineral  weighing  1.0245  g.  yields  0.2602  g.  of 
potassium  and  sodium  chlorides;  and  the  mixed  chlorides  yield  0.4304  g. 
of  K2PtCl6.     Calculate  the  percentage  of  Na20  in  the  sample. 

30.  A  solution  of  chloroplatinic  acid  contains  0.050  g.  of  Pt  per  cubic 
centimeter.  What  is  the  minimum  volume  with  which  0.2602  g.  of  mixed 
sodium  and  potassium  chlorides  must  be  evaporated  in  order  to  insure  the 
complete  conversion  of  the  alkali  metals  into  chloroplatinates,  no  matter 
in  what  proportions  the  two  chlorides  may  exist  in  the  mixture  ? 

31.  A  sample  of  phosphate  rock  contains  0.87%  of  moisture  and  91.92% 
of  calcium  phosphate.  Calculate  the  percentage  of  Ca3(P04)2  which  is 
present  on  the  dry  basis. 

32.  If  2.497  g-  of  a-  fertilizer  containing  4.45%  of  moisture  yields  0.3150 g. 
of  Mg2P207,  what  is  the  percentage  of  P2O6  on  the  dry  basis  ? 

33.  Upon  treatment  with  sulphuric  acid,  1.430  g.  of  a  salt  yields  0.5952  g. 
of  Na2S04  and  101.5  cc.  of  CO2,  measured  moist  at  17°  C.  and  757  mm. 
Calculate  the  percentages  of  Na20  and  CO2  in  the  salt.  (Tension  of  aqueous 
vapor  at  17°=  14.45  mm.) 

34.  A  sample  of  silicate  mineral  weighing  0.8196  g.  yields  0.2082  g. 
of  potassium  and  sodium  chlorides;  and  the  mixed  chlorides  yield  0.1963  g. 
of  potassium  perchlorate.  Calculate  the  percentage  of  K2O  and  of  Na20 
in  the  mineral. 


ANALYTICAL  PROBLEMS  IQS 

35.  What  weight  of  water  is  present  in  one  liter  of  air  which  is  50% 
saturated  with  moisture  at  17°  and  748  mm.  ?     (See  problem  33.) 

36.  1.3250  g.  of  pure  Na2C03  is  dissolved  in  water  and  the  solution 
made  up  accurately  to  250.0  cc.  A  portion  is  carefully  transferred  without 
loss  to  a  platinum  dish  by  means  of  a  pipette  supposed  to  deliver  50.00  cc. 
of  liquid.  After  evaporation  with  hydrochloric  acid,  and  ignition,  the 
sodium  chloride  residue  is  found  to  weigh  0.2927  g.  What  volume  of  this 
solution  does  the  pipette  actually  deliver? 

37.  0.7500  g.  of  a  substance  containing  chlorine  and  bromine  yields 
0.5000  g.  of  Ag(Cl,  Br),  This  mixture  is  heated  in  a  current  of  chlorine, 
which  converts  the  bromide  of  silver  into  the  chloride,  and  the  loss  in  weight 
due  to  this  change  is  found  to  be  0.0683  g-  Calculate  the  percentages  of 
chlorine  and  bromine  in  the  sample. 

38.  From  the  following  data,  calculate  the  percentages  of  chlorine, 
bromine,  and  iodine  in  a  mixture  of  alkali  halides :  Weight  of  sample,  0.1500 
g. ;  weight  of  precipitate  obtained  upon  distilling  the  solution  with  nitrous 
acid  and  converting  the  iodine  into  silver  iodide,  0.1056  g. ;  weight  of  silver 
chloride  and  bromide  from  the  residual  solution,  0.1784  g. ;  weight  of  silver 
chloride  obtained  upon  warming  the  latter  mixture  in  a  current  of  chlorine 
gas,  0.1623  g. 

39.  A  0.5000  g.  sample  of  baking  powder,  known  to  contain  only  NaHCOa 
and  KHC4H4O6,  in  equivalent  proportions,  and  starch,  yields  upon  treat- 
ment with  water  30.5  cc.  of  dry  CO2  (0°  and  760  mm.).  Calculate  the  per- 
centage of  each  salt  in  the  material. 

40.  A  salt  containing  barium,  chlorine,  and  water  of  hydration  gave 
upon  analysis  the  following  data:  Weight  of  sample,  i.oooo  g. ;  weight 
after  heating  (water  driven  off),  0.8522  g. ;  weight  of  silver  chloride  ob- 
tained, 1. 1 735  g. ;  weight  of  barium  sulphate  obtained,  0.9594  g.  Calculate : 
(a)  the  percentage  of  each  constituent ;   (b)  the  formula  of  the  compound. 

41.  If  two  I.oooo  g.  samples  of  a  substance  containing  10.00%  of  MgO 
are  weighed  out,  and  the  precipitate  of  MgNH4P04 .  6  H2O  is  in  one  case 
contaminated  with  0.0250  g.  of  Mg3(P04)2,  and  in  the  other  case  with 
0.0250  g.  of  Mg[(NH4)2P04]2,  what  percentages  of  MgO  will  be  found  if 
the  calculations  are  based  upon  the  assumption  that  the  ignited  precipi- 
tate in  each  case  consists  entirely  of  Mg2P207  ? 

42.  The  carbonates  of  calcium,  strontium,  and  bariimi  obtained  from  a 
lo-liter  sample  of  mineral  water  are  converted  into  the  anhydrous  nitrates, 
and  the  calcium  nitrate  is  extracted  with  absolute  alcohol-ether  mixture. 
The  residue  is  dissolved  in  water,  the  barium  separated  from  the  strontium, 
as  BaCr04,  and  the  strontium  precipitated  from  the  filtrate  with  sulphuric 
acid  and  alcohol.    There  are  finally  obtained  0.8507  g.  of  CaO,  0.1324  g. 


196  QUANTITATIVE  CHEMICAL  ANALYSIS 

of  SrS04,  and  0.1072  g.  of  BaCr04;   calculate  the  content  of  the  water  in 
milligrams  per  liter  {i.e.  parts  per  million)  of  Ca,  of  Sr,  and  of  Ba. 

43.  How  many  cubic  centimeters  of  sulphuric  acid  of  sp.  gr.  1.840,  con- 
taining 95.6%  of  H2SO4,  must  be  added  to  i  liter  of  sulphuric  acid  of  sp. 
gr.  1.560,  containing  65.1%  of  H2SO4,  to  yield  a  solution  containing  75.0% 

0fH2S04? 

44.  A  fuming  sulphuric  acid  contains  25.5%  of  non-hydrated  SO3.  How 
many  grams  of  98.2%  H2SO4  must  be  added  to  100  g.  of  the  fuming  acid  to 
give  a  product  containing  100%  of  H2SO4? 

45.  A  limestone  contains  90.0%  of  CaCOs,  3.50%  of  MgCOa,  3.00%  of 
CaS04  .  2  H2O,  1.25%  of  FeCOa,  and  2.25%  of  anhydrous  siHceous  material. 
What  numerical  difference  would  you  expect  to  find  between  the  loss  on 
ignition  and  the  true  percentage  of  CO2  ? 

46.  An  ore  contains  28.15%  of  nickel,  and  0.5000  g.  samples  are  taken 
for  analysis.  In  one  sample  the  element  is  determined  by  electrolysis,  as 
metallic  nickel,  while  in  a  second  sample  it  is  determined  by  means  of 
dimethylglyoxime,  as  Ni(C4H7N 202)2.  If  the  algebraic  sum  of  the  errors 
involved  in  each  determination  were  equivalent  to  a  negative  error  of 
1.7  mg.  of  the  substance  finally  weighed,  how  much  greater  would  the 
percentage  error  be  in  the  first  determination  than  in  the  second  ? 

47.  An  electric  current  is  passed  simultaneously  through  a  series  of 
three  electrolytic  cells  which  contain  water  acidified  with  sulphuric  acid, 
an  ammoniacal  solution  of  nickel  sulphate,  and  molten  silver  chloride. 
What  is  deposited  upon  the  cathode  in  each  of  the  other  cells,  and  how 
many  grams,  in  the  time  in  which  one  liter  of  hydrogen,  measured  moist 
at  17°  and  746  mm.,  is  liberated  from  the  water?  (Tension  of  aqueous 
vapor  at  17°=  14.45  nim.)  • 

48.  From  the  following  data,  calculate  the  percentages  of  nickel  and 
cobalt  in  the  steel:  Weight  of  sample,  1.124  g. ;  weight  of  nickel  and 
cobalt  obtained  upon  electrolysis,  o.  1 246  g. ;  weight  of  nickel  dimethyl- 
glyoximine,  Ni(C4H7N202)2,  obtained  from  the  electrolytic  deposit,  0.4382  g. 

49.  A  mass  of  platinum  weighs  12.145  g-  i^  air,  11.580  g.  in  water,  and 
I  I.I  15  g.  in  sulphuric  acid.  What  is  the  specific  gravity  of  the  platinum? 
Of  the  sulphuric  acid? 

50.  A  quantity  of  pure  metallic  silver  weighing  1.0788  g.  is  dissolved  in 
nitric  acid  and  the  solution  made  up  to  the  mark  in  a  measuring  flask  gradu- 
ated to  contain  loo.o  cc.  Three  portions  are  carefully  transferred  without 
loss  to  three  separate  beakers  by  means  of  a  pipette  known  to  deliver  25.00  cc. 
If  the  solution  remaining  in  the  flask,  together  with  the  Hquid  finally  washed 
from  the  pipette,  yields  on  analysis  0.3560  g.  of  AgCl,  what  volume  of  liquid 
does  the  flask  actually  contain? 


ANALYTICAL  PROBLEMS  197 

VOLUMETRIC  ANALYSIS 

'^  51.  If  25.00  cc.  of  hydrochloric  acid  yield  0.1435  g.  of  AgCl,  what  is  the 
normality  of  the  solution  ? 

52.  If  a  25.00  cc.  portion  of  acid  requires  21.50  cc.  of  0.526  iV  alkali  for 
neutralization,  what  is  the  normality  of  the  acid?  Supposing  the  acid  to 
be  HCl,  what  weight  of  silver  chloride  will  10.00  cc.  of  it  yield  with  silver 
nitrate  ? 

53.  If  a  2.453  g-  sample  of  pure  anhydrous  sodium  carbonate  requires 
45.72  cc.  of  an  acid  for  neutralization,  and  if  41.90  cc.  of  the  acid 
requires  44.35  cc.  of  an  alkali,  what  is  the  normality  factor  of  each 
solution  ? 

54.  A  sample  of  pure  calcite,  CaCOa,  weighing  2.150  g.  is  dissolved  in 
50.00  cc.  of  an  acid,  and  the  excess  of  acid  is  neutralized  with  29.12  cc.  of 
an  alkali  of  which  28.40  cc.  require  7.10  cc.  of  the  acid  for  neutralization. 
To  what  volume  must  one  liter  of  the  acid  be  diluted  in  order  to  make  it 
exactly  normal? 

>  55.  How  many  cubic  centimeters  of  0.526  N  acid  will  it  take  to  neu- 
tralize the  ammonia  set  free  upon  distilling  1.0378  g.  of  MgNH4P04  .  6  H2O 
with  an  excess  of  caustic  alkali  ? 

56.  If  15.25  cc.  of  alkali  will  neutralize  20.00  cc.  of  a  solution  containing 
6.000  g.  of  KH3(C204)2  •  2  H2O  in  250.0  cc,  what  is  the  normality  factor 
of  the  alkali  ? 
\  57.  In  the  analysis  of  a  feeding  stuff,  a  Kjeldahl  determination  is  carried 
out  with  a  sample  weighing  1.500  g.  The  ammonia  is  received  in  25.00  cc. 
of  0.500  N  acid,  and  the  excess  of  acid  is  found  to  require  12.50  cc.  of  a 
standard  alkali,  of  which  21.20  cc.  will  neutralize  18.00  cc.  of  the  acid. 
Calculate  the  percentage  of  nitrogen  in  the  sample. 

58.  A  sample  of  soda  weighing  25.00  g.  is  dissolved  in  water  and  made 
up  to  250.0  cc,  and  one  fifth  of  this  solution  is  taken  for  titration.  What 
must  be  the  normality  of  the  standard  acid  (assuming  the  alkalinity 
to  be  due  wholly  to  NaaCOs)  in  order  that  twice  the  number  of  cubic 
centimeters  of  acid  used  shall  indicate  the  percentage  of  Na2C03  in  the 
sample  ? 

59.  What  weight  of  argol  (crude  cream  of  tartar)  must  be  taken  for 
titration  in  order  that  each  cubic  centimeter  of  0.2000  N  alkali  used  shall 
represent  2.00%  of  KHC4H4O6? 

60.  A  sample  of  caustic  soda  weighing  4.000  g.  is  dissolved  in  water  and 
made  up  to  one  liter.  A  loo.o  cc.  portion  of  this  solution  requires  for 
neutralization  47.50  cc  of  0.2000  N  acid.  A  second  loo.o  cc.  portion, 
after  treatment  with  barium  chloride  in  slight  excess,  is  diluted  to  200.0  cc. 


"^ 


198  QUANTITATIVE  CHEMICAL  ANALYSIS 

and  allowed  to  settle,  and  50.0  cc.  of  the  clear  solution  require  11.50  cc.  of 
the  acid.  Calculate  the  percentages  of  NaOH  and  Na2C03  in  the  sample. 
(Nedect  the  volume  occupied  by  the  solid  precipitate.) 

(6r;  A  sample  ofrS"6lvay  soda  weighing  3.750  g.  is  dissolved  in  water  and 
ina^e  up  to  one  liter.  A  loo.o  cc.  portion  of  this  solution,  titrated  in  the 
cold  with  o.iooo  N  acid,  with  the  use  of  phenolphthalein,  is  found  to  re- 
quire 29.95  cc.  of  the  acid^  the  burette  is  then  refilled  and  the  titration 
completed  at  the  boiling  temperature  of  the  solution,  35.15  cc.  more  of 

_j;he  acid  being  required.  Calculate  the  percentages  of  Na2C03  and  NaHCOs 
in  the  sample.  (Under  suitable  experimental  conditions,  phenolphthalein 
becomes  colorless  in  the  cold  as  soon  as  the  carbonate  has  been  wholly 
converted  into  bicarbonate.) 

y  62.  A  sample  of  sirupy  phosphoric  acid  weighing  5.767  g.  is  dissolved  in 
watef'and  made  up  to  one  liter.  A  ioo;o  cc.  portion  of  the  solution  is  treated 
with  sodium  acetate  and  silver  nitrate  in  excess,  whereby  the  phosphate  is 
quantitatively  precipitated  as  Ag3P04.  Phenolphthalein  is  added  to  the 
filtrate  and  washings,  and  the  solution  titrated  with  0.500  N  alkali,  of  which 
27.25  ca  are  required.  Calculate  the  percentage  of  H3PO4  in  the  original 
sample.  -^^n 

/  63.  A  sample  of  Chili  saltpeter  weighing/ 1.025  g.  is  treated  in  sodium 
hydroxide  solution  with  p«lverized-Devarda's  alloy  (50%  Cu,  45%  Al, 
.5%,Zn),  which  reduces  the  nitrogen  to  ammonia;  the  ammonia  is  distilled 
into  25.00  cc.  of  0.463  N  acid,  and  the  excess  of  acid  requires  5.01  cc.  of 
0.212  N  alkaU  for  neutralization.  Assuming  that  nitrogen  was  wholly 
present  as  NaNOs,  calculate  the  percentage  of  the  latter  in  the  sample. 

64.  A  sample  of  strontium  nitrate  weighing  10.53  g-  is  dissolved  in  water 
and  made  up  to  one  liter.  One  tenth  of  this  solution  is  distilled,  in  the 
presence  of  alkali,  with  an  excess  of  titanous  chloride,  which  reduces 
the  nitrate  to  ammonia  (KN03+8Ti(OH)3+6  H20=KOH+8Ti(OH)4^- 
NH3).  The  ammonia  is  received  in  25.00  cc.  of  0.500  N  acid,  and  the  ex- 
cess of  acid  requires  10.16  cc.  of  0.250  N  alkali.  Assuming  that  the  nitro- 
gen was  wholly  present  as  Sr(NOs)  2,  calculate  the  percentage  of  the  latter  in 
the  sample. 

65.  How  many  grams  of  KH3(C204)2 .  2  H2O  will  it  take  to  prepare  one 
liter  of  a  0.500  N  solution,  to  be  used  as  a  standard  acid?  How  many  to 
prepare  one  liter  of  a  o.iooo  N  solution,  to  be  used  as  a  reducing  agent  in 
connection  with  potassium  permanganate? 

66.  How  many  grams  of  K2Cr207  per  liter  will  be  required  to  prepare  a 
solution  of  such  strength  that  each  cubic  centimeter  shall  indicate  2.00% 
of  iron,  when  a  sample  weighing  0.2792  g.  is  used  for  analysis?  What 
is  the  normality  factor  of  this  solution? 


ANALYTICAL  PROBLEMS  199 

67.  From  the  following  data,  calculate  the  percentage  of  iron  in  the 
ore :  Weight  of  sample,  0.2186  g. ;  the  reduced  iron  solution  requires  for 
oxidation  25.14  cc.  of  0.0996  N  permanganate  solution. 

68.  What  is  the  maximum  weight  of  an  ore  containing  70.00%  of  iron 
which  can  be  taken  for  analysis  without  having  to  refill  a  30-cc.  burette, 
if  the  permanganate  solution  is  0.1025  N? 

69.  A  calcium  oxalate  precipitate,  obtained  from  0.8432  g.  of  a  rock, 
is  decomposed  with  dilute  sulphuric  acid  and  made  up  to  250.0  cc. ;  of  this 
a  50.0  cc.  portion  is  titrated  with  0.1012  N  permanganate  solution.  If 
27.35  cc.  of  the  latter  are  required,  what  percentage  of  CaO  does  the  rock 
contain  ? 

70^^  11,56  cc.  of  nitric  acid  of  sp.  gr.  1.19  are  diluted  to  250.0  cc. ;  20.00 
cc.  (rf 'this  solution  are  found  to  require  12.92  cc.  of  0.410  iV  alkali.  Cal- 
culate the  percentage  of  HNO3  in  the  acid  of  sp.  gr.  1.19.  '    ' 

71.  The  Sb2S3  precipitate  obtained  from  the  solution  of  an  ore  is  dis- 
solved in  sodium  sulphide  solution,  and  this  is  evaporated  and  fumed  with 
an  excess  of  sulphuric  acid.  The  residue  is  then  dissolved  in  dilute  hydro- 
chloric acid,  and  the  antimony  oxidized  from  the  trivaJent  to  the  penta 
valent  condition  by  means  of  a  standard  solution  of  permanganate.  The 
sample  of  ore  weighs  0.2749  g.  and  24.17  cc.  of  0.1025  N  permanganate 
solution  are  used  in  the  titration ;  what  is  the  percentage  of  antimony  ? 

72.  What  is  the  normality  factor  of  an  acid,  of  which  25.37  cc.  are  equiva- 
lent to  1.263  g-  of  KNO3  when  the  nitrogen  of  the  latter  is  reduced  in  alkaline 
solution  to  ammonia  and  this  is  distilled  off  and  received  in  the  acid  solution? 

73.  From  the  following  data,  what  is  the  percentage  of  Mn02  in  the  ore? 
Weight  of  sample,  0.2000  g. ;  after  heating  this  in  the  presence  of  sulphuric 
acid  with  50.00  cc.  of  o.iooo  N  oxalic  acid,  the  excess  of  oxalic  acid  re- 
quires 8.50  cc.  of  1. 0000  N  permanganate  solution. 

74.  From  the  following  data,  what  is  the  percentage  of  Mn02  in  the  ore  ? 
Weight  of  sample,  0.2500  g. ;  this  is  boiled  with  hydrochloric  acid  and  the 
distillate  received  in  an  excess  of  potassium  iodide  solution,  the  liberated 
iodine  requiring  25.51  cc.  of  0.2000  N  thiosulphate  solution. 

75.  A  standard  solution  of  permanganate  will  oxidize  5.783  mg.  of 
ferrous  iron  per  cubic  centimeter;  to  what  volume  must  one  liter  of  the 
solution  be  diluted  in  order  to  obtain  an  exactly  tenth-normal  solution  ? 

76.  A  sample  of  mineral  substance  weighing  i.ooo  g.  is  taken  for  analysis. 
In  the  determination  of  the  iron,  the  ferric  solution  is  completely  reduced 
by  means  of  sulphurous  acid,  and  the  excess  of  the  latter  removed  by  the 
passage  of  carbon  dioxide  through  the  boiling  solution;  the  iron  then  re- 
quires 28,17  cc.  of  o.iooo  N  permanganate  solution.  Calculate  the  per- 
centage of  iron  in  the  substance. 


200  QUANTITATIVE  CHEMICAL  ANALYSIS 

77.  What  weight  of  Fe(NH4S04)2  •  6H2O  will  reduce  50.0  cc.  of  a 
permanganate  solution,  of  which  10.00  cc.  will  liberate  from  an  acidified 
solution  of  potassium  iodide  a  quantity  of  iodine  sufficient  to  react  with 
15.25  cc.  of  0.1025  N  thiosulphate  solution? 

78.  The  calcium  oxalate  precipitate  obtained  from  0.2500  g.  of  calcite 
is  dissolved  in  an  excess  of  sulphuric  acid,  and  the  hot  solution  titrated  with 
a  solution  of  permanganate  of  which  each  cubic  centimeter  represents 
0.00735  g-  of  Na2C204.  If  45.57  cc.  of  the  permanganate  solution  are  re- 
quired, what  percentage  of  calcium  does  the  mineral  contain  ? 

79.  A  standard  solution  of  permanganate  will  oxidize  0.00730  g.  of  ferrous 
iron  per  cubic  centimeter ;  what  is  the  value  of  the  same  solution  in  terms  of 
{a)  H2C2O4 .  2  H2O;  (6)  NaN02;  {c)  H2O2;  {d)  K4Fe(CN)6 .  3  H2O;  {e)  Mn? 

80.  If  0.1340  g.  of  sodium  oxalate  require  19.23  cc.  of  a  permanganate 
solution,  how  many  milligrams  of  ferrous  iron  will  each  cubic  centimeter 
of  the  permanganate  solution  indicate  ? 

81.  From  the  following  data,  calculate  the  percentage  of  manganese 
in  the  ore:   Weight  of  sample,  0.5027  g. ;    volume  of  permanganate  solu 
tion  required  to  oxidize  the  manganese,  36.60  cc. ;    value  of  the  perman- 
ganate solution  for  use  with  oxalic  acid,  0.0997  ■^• 

82.  2.400  liters  of  dry  air  (at  0°,  760  mm.)  and  50.00  cc.  of  o.oioo  N 
barium  hydroxide  solution  are  shaken  together,  and  the  excess  of  alkali 
is  found  to  require  35.06  cc.  of  0.0100  N  acid.  What  volume  of  carbon 
dioxide  is  contained  in  10,000  volumes  of  the  dry  air? 

83.  If  a  permanganate  solution  is  equivalent  to  5.84  mg.  of  iron  per  cubic 
centimeter,  what  is  the  value  of  the  solution  in  terms  of  K4Fe(CN)6 .  3  H2O? 
Intermsof  N2O3? 

84.  What  weight  of  iodine  per  cubic  centimeter  will  be  liberated  by  a 
permanganate  solution  from  an  excess  of  hydriodic  acid,  if  the  perman- 
ganate solution  has  an  iron  value  of  4.98  mg.  per  cubic  centimeter? 

85.  A  piece  of  iron  wire  weighing  0.1408  g.,  and  containing  99.84% 
of  iron,  is  converted  into  ferrous  chloride  and  titrated  according  to  the  Zim- 
mermann-Reinhardt  method;  it  requires  25.15  cc.  of  a  permanganate  solu- 
tion. What  weight  of  ore  must  be  taken  for  analysis  by  the  same  method, 
in  order  that  each  cubic  centimeter  of  the  permanganate  solution  shall  in- 
dicate 2.50%  of  iron? 

86.  25.00  cc.  of  a  certain  acid  are  found  to  require  23.67  cc.  of  an  alkaH. 
If  28.15  cc.  of  the  acid  are  used  to  dissolve  0.5260  g.  of  pure  calcium  car- 
bonate, and  6.61  cc.  of  the  alkali  are  required  to  neutralize  the  resulting 
solution,  what  is  the  normality  factor  of  the  acid?  What  is  that  of  the 
alkali? 

87.  The  calcium  oxalate  from  0.5165  g.  of  a  mineral  requires,  after 


ANALYTICAL  PROBLEMS  201 

decomposition  with  sulphuric  acid,  43.56  cc.  of  a  permanganate  solution 
having  a  sodium  oxalate  value  of)  6. 70  mg.  per  cubic  centimeter.  What 
is  the  percentage  of  CaO  in  the  mineral  ? 

88.  In  the  standardization  of  a  dichromate  solution,  a  sample  of  pure 
iron  weighing  0.2000  g.  is  converted  into  200  cc.  of  ferrous  chloride  solution, 
and  titrated.  In  the  subsequent  analysis  of  an  ore  with  the  dichromate 
solution,  an  equal  quantity  of  ferrous  iron  is  present,  but  the  solution  to  be 
titrated  has  a  volume  of  600  cc.  If  the  indicator  used  permits  the  recog- 
nition of  one  part  by  weight  of  ferrous  iron  in  200,000  of  solution,  what  error 
results  from  the  fact  that  the  two  titrations  are  made  at  different  volumes  ? 
(Assume  that  the  dichromate  solution  was  found  to  be  o.iooo  N,  and  that 
the  solutions  have  the  specific  gravity  of  water.) 

89.  A  sample  of  crystalline  ammonium  acetate  weighing  2.021  g.  is 
dissolved  in  water  and  made  up  to  200.0  cc.  One  half  of  this  solution  is 
distilled  with  an  excess  of  lime,  and  the  distillate  received  in  25.00  cc.  of 
0.500  N  acid;  the  second  half  is  distilled  with  phosphoric  acid  in  excess, 
and  the  distillate  received  in  45.00  cc.  of  0.500  N  alkali.  In  the  first  case, 
with  methyl  orange,  the  excess  of  acid  requires  10.20  cc.  of  the  standard 
alkali,  and  in  the  second  case,  with  phenolphthalein,  the  excess  of  alkali 
requires  15.52  cc.  of  the  standard  acid.     Calculate  the  formula  of  the  salt. 

90.  The  aqueous  solution  of  o.  1361  g.  of  a  mixture  containing  only  sodium 
chloride  and  bromide  is  treated,  in  the  presence  of  nitric  acid,  with  25.00  cc. 
of  O.IOOO  N  AgNOa  solution,  and  the  precipitate  is  filtered  off  and  washed. 
The  filtrate  and  washings  require  9.70  cc.  of  o.iooo  N  thiocyanate  solution. 
Calculate  the  percentages  of  sodium  chloride  and  bromide  in  the  sample. 

91.  A  mixture  containing  soluble  chlorides  and  iodides  weighs  0.4500  g. 
This  is  treated,  in  the  presence  of  nitric  acid,  with  35.00  cc.  of  o.iooo  N 
silver  nitrate  solution,  and  the  precipitate  is  found  to  weigh  0.5000  g.  The 
excess  of  silver  nitrate  in  the  filtrate  and  washings  requires  11. 10  cc.  of 
0.0500  N  thiocyanate  solution.  Calculate  the  percentages  of  chlorine  and 
iodine  in  the  original  mixture. 

92.  A  20.00  cc.  portion  of  a  pnissic  acid  solution  is  made  slightly  alka- 
line with  sodium  hydroxide,  and  a  very  little  potassium  iodide  is  added; 
the  solution  is  then  titrated  with  o.iooo  iV  silver  nitrate,  of  which  48.73 
cc.  are  required  to  produce  a  faint  permanent  turbidity  of  silver  iodide. 
Calculate  the  percentage  of  HCN  in  the  original  solution,  assuming  the 
specific  gravity  of  the  latter  to  be  that  of  water. 

93.  A  mixture  containing  potassium  cyanide  and  chloride,  and  weigh- 
ing 0.2037  g.,  is  dissolved  in  water  and  titrated  with  o.iooo  N  silver  ni- 
trate solution,  of  which  14.41  cc.  are  required  to  produce  a  faint  permanent 
turbidity.     25.00  cc.  more  of  the  silver  nitrate  solution  are  added,  the 


202  QUANTITATIVE  CHEMICAL  ANALYSIS 

solution  is  slightly  acidified  with  nitric  acid,  and  the  filtrate  and  washings 
from  the  precipitate  require  for  titration  8.91  cc.  of  o.iooo  N  thiocyanate 
solution.     Calculate  the  percentages  of  KCN  and  KCl  in  the  sample. 

94.  A  solution  of  potassium  permanganate  is  equivalent  to  5.00  mg.  of 
iron  per  cubic  centimeter.  To  40.0  cc.  of  this  solution,  acidified  with  an 
excess  of  very  dilute  sulphuric  acid,  an  excess  of  potassium  iodide  is  added 
and  the  Hberated  iodine  is  titrated  with  a  solution  of  sodium  thiosulphate, 
of  which  34.85  cc.  are  required.  Calculate  the  normality  factor  of  the 
thiosulphate  solution. 

95.  A  sample  of  bleaching  powder  weighing  7.092  g.  is  triturated  with 
water  and  made  up  to  one  liter.  A  50.0  cc.  portion  of  this  suspension, 
when  titrated  with  o.iooo  N  arsenious  acid,  with  potassium  iodide-starch 
paper  as  an  outside  indicator,  is  found  to  require  26.15  cc.  of  the  standard 
solution.  What  is  the  percentage  of  available  chlorine  in  the  bleaching 
powder? 

96.  A  sample  of  soda  weighing  4.973  g.  is  dissolved  in  water  and  made 
up  to  one  liter.  A  loo.o  cc.  portion  of  this  solution,  carefully  titrated  in 
the  cold  with  phenolphthalein  as  an  indicator,  requires  48.90  cc.  of  0,0998  N 
acid.  A  second  portion  of  loo.o  cc.  is  titrated  with  0.499  N  acid,  of  which 
13.81  cc.  are  required  with  methyl  orange  as  an  indicator.  Calculate  the 
percentages  of  Na2C03  and  NaOH  in  the  sample.     (Cf.  Problem  61.) 

97.  If  43.60  cc.  of  a  thiosulphate  solution  require  40.15  cc.  of  an  iodine 
solution,  and  if  0.2 118  g.  of  AS2O3  require  42.40  cc.  of  the  iodine  solution, 
what  is  the  normality  factor  of  each  solution? 

98.  A  sample  of  titaniferous  ore  weighing  0.3805  g.  is  fused  with  a  mix- 
ture of  K2S2O7  and  KF,  the  melt  dissolved  in  HCl,  and  the  iron  and  titanium 
reduced  with  zinc  in  an  atmosphere  of  hydrogen.  The  solution  is  then 
titrated  in  an  atmosphere  of  carbon  dioxide,  in  the  presence  of  i  g.  of  KSCN 
as  the  indicator,  with  o.iooo  N  ferric  chloride  solution,  of  which  19.34  cc. 
are  required.  Calculate  the  percentage  of  Ti02  in  the  ore.  (TiCls+FeCla 
=  TiCl4+FeCl2.) 

99.  In  50.0  cc.  of  a  solution,  containing  both  ferrous  and  ferric  sulphates, 
the  ferrous  iron  is  titrated  in  the  presence  of  sulphuric  acid  with  o.iooo  N 
permanganate  solution,  after  which  the  oxidized  solution  is  titrated  for  total 
iron  with*  0.0997  ^  titanous  chloride  solution,  with  potassium  thiocyanate 
as  the  indicator.  If  in  the  first  titration  27.17  cc.  of  the  permanganate 
solution  are  used,  and  in  the  second  titration  46.98  cc.  of  the  titanous 
chloride  solution,  what  is  the  content  of  the  original  solution  in  grams  per 
liter  of  ferrous  and  of  ferric  iron? 

100.  The  ammonium  phosphomolybdate  precipitate  obtained  from 
2.000  g.  of  steel  is  dissolved  in  dilute  aqueous  ammonia,  the  solution  acidified 


ANALYTICAL  PROBLEMS  203 

with  sulphuric  acid,  and  the  molybdenum  reduced  to  the  trivalent  con- 
dition by  passing  the  acid  solution  through  a  Jones  reductor,  —  the  re- 
duced solution  being  caused  to  enter  the  receiving  vessel  under  the  sur- 
face of  a  solution  of  ferric  sulphate  (2  M0O3+6  H=Mo203-h3  H2O;  and 
M02O3+3  Fe203=2  M0O3+6  FeO).  The  resulting  solution  is  at  once 
titrated  with  a  standard  solution  of  permanganate,  of  which  18.75  cc.  are 
required.  If  the  permanganate  solution  has  an  iron  value  of  0.00540  g,  per 
cubic  centimeter,  what  is  the  percentage  of  phosphorus  in  the  steel? 


APPENDIX 

THE  PREPARATION  OF  THE  REAGENTS 

Many  of  the  reagents  which  are  used  in  quantitative  analysis 
are  made  up  for  one  purpose  only ;  directions  for  the  preparation 
of  these  are  given  in  this  book  under  the  individual  determina- 
tions. Certain  reagents,  however,  are  used  in  many  operations 
at  approximately  fixed  concentrations,  and  it  is  especially  these 
which  are  included  in  this  section. 

It  is  of  great  advantage  to  have  such  stock  reagents  conform 
to  some  definite  system  of  concentration,  the  most  convenient 
system  being  one  based  upon  the  equivalent  or  normal  weights 
employed  in  volumetric  analysis.  Of  course,  however,  the  con- 
centrations of  such  solutions  need  not  be  so  exactly  fixed  as  for 
use  in  accurate  volumetric  work. 

With  this  system,  equal  volumes  of  the  solutions  bear  fixed 
relations  to  one  another,  it  is  easy  to  calculate  the  volume  of  a 
reagent  required  for  a  specific  purpose,  and  the  addition  of  an  un- 
necessary excess  may  readily  be  avoided.  Thus,  if  we  dissolve 
I  g.  of  calcium  carbonate  in  lo  cc.  of  6-normal  hydrochloric  acid, 
and  wish  to  neutralize  the  liquid  with  ammonia,  we  see  at  once  that 

=  6.7  cc.  will  be  about  the 


10  — 


CaC03X6 


2000 


10  — 


100X6 


2000 

right  amount.  Again,  if  we  fuse  0.15  g.  of  chromite  with  4.0  g. 
of  Na202,  extract  the  fusion  with  water,  and  wish  to  almost 
(but  not  quite)  neutralize  the  aqueous  extract,  we  see  that 

— — ^ 7=    ^'    ^  =  17.1  cc.  of  6-normal  acid  would  neutralize 

Na202X6      78X6       ' 

2000  2000 

20s 


206  QUANTITATIVE  CHEMICAL  ANALYSIS 

the  liquid;  consequently  we  use  15-16  cc.  of  the  acid,  and  test 
the  resulting  mixture  to  make  sure  of  its  alkalinity.  As  will 
be  realized,  this  e£fects  a  saving  in  time,  labor,  and  material; 
and  also  it  leads  to  more  accurate  and  reliable  work. 

Graduated  cylinders  and  measuring  pipettes  are  useful  for 
delivering  specific  volumes  of  such  reagents. 

ACIDS 

Acetic,  6-normal:  Mix  350  cc.  of  glacial  acetic  acid  with 
650  cc.  of  water. 

Hydrochloric y  12-normal:  Use  the  C.  P.  acid  of  commerce 
of  sp.  gr.,  1. 19. 

Hydrochloric,  6-normal:  Mix  12-normal  acid  with  an  equal 
volume  of  water.  The  specific  gravity  of  this  acid  is  about 
1. 10. 

Nitric,  1 6-normal:  Use  the  C.  P.  acid  of  commerce  of  sp.  gr., 
1.42. 

Nitric,  6-normal:  Mix  380  cc.  of  the  1 6-normal  acid  with 
650  cc.  of  water.  The  specific  gravity  of  this  acid  is  about 
1.195. 

Sulphuric,  3 6-normal :  Use  the  C.  P.  acid  of  commerce  of  sp. 
gr.,  1.84. 

Sulphuric,  6-normal :  Pour  200  cc.  of  the  3 6-normal  acid 
into  1045  cc.  of  water.     The  specific  gravity  of  this  acid  is  1.18. 

BASES 

Ammonium  hydroxide,  15-normal:  Use  the  C.  P.  ammonia 
water  of  commerce  of  sp.  gr.,  0.90. 

Ammonium  hydroxide,  6-normal:  Mix  400  cc.  of  the  15- 
normal  solution  with  600  cc.  of  water.  The  specific  gravity 
of  this  solution  is  about  0.958. 

Sodium  hydroxide,  6-normal :  Dissolve  250  g.  of  stick  sodium 
hydroxide  in  water  and  dilute  to  one  liter. 


APPENDIX  307 


SALTS 


Ammonium  carbonate:  Dissolve  250  g.  of  freshly  powdered 
ammonium  carbonate  in  one  liter  of  6-normal  ammonium  hy- 
droxide, and  filter  if  there  is  a  residue. 

Ammonium  molyhdate:  ^  Dissolve  100  g.  of  M0O3  in  80  cc. 
of  ammonia  (sp.  gr.,  0.90)  with  the  addition  of  400  cc.  of  water ; 
with  cooling  and  constant  stirring,  allow  the  clear  solution  to 
run  slowly  into  a  mixture  of  400  cc.  of  nitric  acid  (sp.  gr., 
1.42)  with  600  cc.  of  water,  add  0.05  g.  of  microcosmic  salt, 
NaNH4HP04 .  4  H2O,  and  keep  the  mixture  in  a  warm  place  for 
several  days,  or  until  a  portion  heated  to  40°  deposits  no  yellow 
precipitate.  Decant  from  any  sediment,  and  preserve  in  glass- 
stoppered  bottles.     This  solution  contains  68  g.  of  M0O3  per  liter. 

Ammonium  oxalate,  0.5-normal:  Dissolve  35.5  g.  of 
(NH4)2C204  .  H2O  in  1000  cc.  of  water. 

Barium  chloride,  i -normal:  Dissolve  122  g.  of  BaCU  .  2  H2O 
in  1000  cc.  of  water. 

Magnesia  mixture,  0.5-normal  as  a  precipitant  for  phosphoric 
or  arsenic  acid:  Dissolve  51  g.  of  MgCl2 .  6  H2O  and  130  g. 
of  NH4CI  in  water,  add  121  cc.  of  ammonia  (sp.  gr.,  0.90),  and 
dilute  to  one  liter. 

Mercuric  chloride,  0.2-normal  for  oxidizing  stannous  chloride: 
Dissolve  54  g.  of  HgCl2  in  1000  cc.  of  hot  water. 

Silver  nitrate,  0.2-normal :  Dissolve  34  g.  of  AgNOs  in  1000  cc. 
of  water. 

Sodium  phosphate,  0.5-normal  as  a  precipitant  for  mag- 
nesium: Dissolve  90  g.  of  Na2HP04 .  12  H2O  (or  52  g.  of 
NaNH4HP04 .  4  H2O)  in  1000  cc.  of  water. 

1  Recovery  of  the  Molybdic  Acid.  To  the  liquid  molybdate  residues,  acidified 
if  necessary  with  nitric  acid,  add  sodium  phosphate  solution  in  excess.  Collect 
the  yellow  precipitate,  wash  it  with  water  containing  sodium  sulphate,  and  then 
dry  it  in  the  air.  Dissolve  i  Kg.  of  the  dried  precipitate  in  ammonia,  add  a  strong 
solution  of  60  g.  of  NH4CI  and  1 20  g.  of  MgCU  .  6  H2O  in  water,  allow  to  stand  for 
6  hours,  and  filter  off  the  precipitate.  To  the  filtrate,  decolorized  if  necessary  with 
a  little  H2O2,  add  HCl  just  to  acid  reaction,  to  precipitate  the  M0O3.  Collect 
this  precipitate,  wash  with  water,  and  dry  at  1 10°. 


2o8  QUANTITATIVE  CHEMICAL  ANALYSIS 

Stannous  chloride,  i -normal  as  a  reducing  agent:  Dissolve 
113  g.  of  SnCU  .  2  H2O  in  150  cc.  of  12-nonnal  hydrochloric 
acid,  with  the  gradual  addition  of  water,  finally  diluting  to  one 
liter.     Keep  in  bottles  containing  granulated  tin. 

Sulphuric  Acid- Dichr ornate  Cleaning  Solution 

With  stirring,  cautiously  pour  200  cc.  of  sulphuric  acid  (sp. 
gr.,  1.84)  into  150  cc.  of  cold  water,  and  saturate  the  hot  solu- 
tion, without  further  heating,  with  powdered  sodium  (or  potas- 
sium) dichromate. 

When  cleaning  measuring  vessels  with  this  liquid,  they  should 
be  filled  with  the  cold  solution  and  allowed  to  stand  overnight, 
or  longer. 

Analytical  Samples  for  the  Use  of  Students 

The  analyzed  samples  indicated  in  the  text  for  the  use  of 
beginners  in  quantitative  analysis  may  in  some  cases  be  ob- 
tained in  the  market.  Otherwise,  they  may  be  prepared  by 
mixing  together  the  component  materials  in  the  proportions 
decided  upon.  This  mixing  is  best  accomplished  by  long  con- 
tinued grinding  in  a  ball  mill,  the  material  being  finally  passed 
through  a  fine-meshed  sieve,  and  bottled.  These  samples  should 
be  carefully  analyzed  by  members  of  the  quantitative  staff,  so 
that  the  student's  work  may  be  judged  according  to  its  accuracy. 
Most  of  the  mixtures  indicated  can  be  kept  from  year  to  year 
without  change. 

It  is  desirable  to  have  in  each  case  a  continuous  series  of 
at  least  ten  samples,  varying  in  content  from  sample  to  sample 
by  about  0.4-0.5%. 


APPENDIX 


209 


APPARATUS    IN  THE  STUDENT'S  DESK^ 
Quantitative  Chemical  Laboratories 


Above  in  the  drawers 
I  Brush,  camel's  hair. 
I  Burette,  g.  s.,  30  cc. 

1  Burette,  30  cc,  for  pinchcock. 
4  Crucibles,  porcelain,  o. 

2  Crucibles,  porcelain,  Gooch,  extra 

disc. 
2  Cylinders,  graduated,  50  cc.  and 

10  cc. 
I  File. 

1  Forceps,  steel. 

2  Funnels,    diam.     25    mm.,    stem 

40  mm. 
2  Glasses,  watch,  140  mm. 
2  Glasses,  watch,  70  mm. 
2  Glasses,  watch,  50  mm. 
I  Vial  litmus  paper,  blue. 

1  Vial  litmus  paper,  red. 

2  Boxes  matches,  safety. 
I  Pinchcock. 

I  Pipette,  25  cc. 

1  Pipette,  10  cc. 

2  Policemen,  rubber  tip. 

3  Rods,  glass,  200  mm. 
6  Test  tubes. 

I  Thermometer,  100°  C. 

I  Tongs,  brass,  nickel  plated. 

1  Tube,  connecting,  3 -way. 

2  Tubes,  rubber,  for  Gooch  crucibles. 

3  Tubes,  weighing,  with  corks. 

I  Tube,     rubber,    pressure,    length 

300  mm. 
I  Tube,      rubber,      small,      length 

300  nam. 
Tubing,  soft  glass,  900  mm. 


Below  in  the  cupboard 
12  Beakers,  2  nests,  1-6,  with  win- 
dow pane. 
2  Bottles,  g-  s.,  2500  cc. 
I  Bottle,  g.  s.,  250  cc,  for  cleaning 

solution. 
I  Bottle,  g.  s.,  125    cc,  for  silver 

nitrate. 
I  Bottle,  weighing. 

1  Burette  holder,  Lincoln's 

2  Burners,  adjustable. 

2  Burner  tubes,  rubber. 

2  Casseroles,  porcelain,  500  cc. 

1  Desiccator  for  4  crucibles. 

2  Flasks,  Erlenmeyer,  700  cc. 
2  Flasks,  Erlenmeyer,  500  cc. 
2  Flasks,  Erlenmeyer,  250  cc. 
2  Flasks,  Erlenmeyer,  150  cc. 
2  Flasks,  filter,  500  cc 

2  Flasks,  Florence,  500  cc. 
2  Flasks,  Florence,  250  cc. 
2  Flasks,  Florence,  50  cc,  for  in- 
dicators. 
I  Flask,  volumetric,  1000  cc. 

1  Flask,  volumetric,  500  cc 

2  Flasks,  volumetric,  250  cc. 

1  Flask,  volumetric,  100  cc. 

4  Funnels,    diam.    70    mm.,    stem 
200  mm. 

2  Funnels,  for  Gooch  crucibles. 
I  Sponge. 

1  Stand,  filter,  wooden. 

2  Stands,  iron,  i  ring  each. 

2  Triangles,  pipe  stem,  new  form. 
2  Tripods,  iron. 


2  Wire  gauzes. 
*  The  articles  listed  above  represent  the  apparatus  with  which  it  is  desirable 
to  provide  each  student  at  the  outset ;   the  list  can  of  course  be  modified  in  many 
particulars  without  jeopardizing  the  success  of  the  work.     Any  additional  appa- 
ratus which  may  be  required  can  be  obtained  as  needed  from  the  storeroom. 


2IO 


Table  A — Four  Place  Logaritliiiis 


[A 


IT 

0 

1 

2 

3 

4 

5 

6 

7 

8 

9 

12  3 

4  5  6 

7  8  9 

10 

0000 

0043 

0086 

0128 

0170 

0212 

0253 

0294 

0334 

0374 

4  8  12 

17  21  25 

29  33  37 

11 
12 
13 

14 
15 
16 

17 
18 
19 

0414 
0792 
1139 

1461 
1761 
2041 

2304 

2553 

2788 

0453 
0828 
1173 

1492 

1790 
2068 

2330 

2577 
2810 

0492 

0864 
1206 

1523 
1818 
2095 

2355 
2801 
2833 

0531 

0899 
1239 

1553 

1847 
2122 

2380 
2625 
2856 

056» 
0934 
1271 

1584 
1875 
2148 

2405 
2648 
2878 

0607 
0969 
1303 

1614 
1903 
2175 

2430 
2672 
2900 

0645 
1004 
1335 

1644 
1931 
2201 

2455 
2695 
2923 

0682 
1038 
1367 

1673 
1959 

2227 

2480 
2718 
2945 

0719 
1072 
1399 

1703 
1987 
2253 

2504 

2742 
2967 

0755 
1106 
1430 

1732 
2014 
2279 

2529 
2765 
2989 

4  8  11 
3  7  19 
3  610 

3  0  9 
3  6  8 
3  5  8 

2  5  7 
2  5  7 
2  4  7 

15  19  23 
14  17  21 
131619 

121518 
11  14  17 
11 13  16 

10  12  15 
9  1214 
9  1113 

26  30  34 
24  28  31 
23  26  29 

2124  27 

20  22  25. 
18  21  24 

17  20  22 
16  19  21 

16  18  20 

20 

3010 

3032 

3054 

3075 

3096 

3118 

3139 

3160 

3181 

3201 

2  4  6 

8  1113 

15  17  19 

21 
22 
23 

24 
25 

26 

27 

28 
29 

3222 
3424 
3617 

3802 
3979 
4150 

4314 
4472 
4624 

3243 
3444 
3636 

3820 
3997 
4166 

4330 

4487 
4639 

3263 
3464 
3655 

3838 
4014 
4183 

4346 

4502 
4654 

3284 
3483 
3674 

3856 
4031 
4200 

4362 
4518 
4069 

3304 
3502 
3692 

3874 
4048 
4216 

4378 
4533 
4683 

3324 
3522 
3711 

3892 
4065 
4232 

4393 

4548 
4698 

3345 
3541 
3729 

3909 
4082 
4249 

4409 
4564 
4713 

3365 
3560 
3747 

3927 
4099 
4265 

4425 
4579 
4728 

3385 
3579 
3766 

39i5 
4116 
4281 

4440 

4594 
4742 

3404 
3598 
3784 

3962 
4133 
4298 

4456 
4609 
4757 

2  4  6 
2  4  6 
2  4  6. 

2  4  6 
2  4  5 
2  3  5 

2  3  5 
2  3  5 
13  4 

8  1012 
8  1012 
7  9  11 

7  9  11 
7  9  10 
7  8  10 

6  8  9 
6  8  9 
6  7  9 

14  16  18 
14  16  17 
13  15  17 

121416 
12  14  10 
11 13  15 

11 12  14 
11 12  14 
10  12  13 

30 

4771 

4786 

4800 

4814 

4829 

4843 

4857 

4871 

4886 

4900 

13  4 

6  7  9 

10  11 13 

31 
32 
33 

34 
35 

36 

37 

38 
39 

4914 
5051 
5185 

5315 
5441 
5563 

5682 
5798 
5911 

4928 
5065 
5198 

5328 
5453 
5575 

5694 
5809 
5922 

4942 
5079 
5211 

5340 
5405 
5587 

5705 
5821 
5933 

4955 
5092 
5224 

5353 

5478 
5599 

5717 
5832 
5944 

4969 
5105 
5237 

5366 
5490 
5611 

5729 
5843 
5955 

4983 
5119 
5250 

5378 
5502 
5623 

5740 
5855 
5966 

4997 
5132 
6263 

5391 
5514 
5635 

5752 

58G6 
5977 

5011 
5145 
5276 

5403 
5527 
5647 

5763 
5877 
50C3 

5024 
6159 
6289 

5416 

5539 
5658 

57V5 
6888 
5999 

5038 
5172 
6302 

5428 
5551 
5670 

5786 
5899 
6010 

13  4 
13  4 
13  4 

12  4 

12  4 
12  4 

12  4 

12  3 
1  2  3 

6  7  8 

5  7  8 

6  7  8 

5  6  8 

5  6  7 

6  6  7, 

6  6  7 
6  6  7 

4  6  7 

10  11 12 
9U12 
91112v 

91011 
9  10  11 
81011 

8  911 
8  910 
8  910 

40 

6021 

6031 

6042 

6053 

6064 

6075 

6085 

0096 

6107 

6117 

12  3 

4  6  6 

8  910 

41 

42 
43 

44 
45 

46 

47 

48 
49 

6128 
6232 
6335 

6435 
6532 
6628 

6721 
6812 
6902 

6138 
6243 
6345 

6444 
6542 
6637 

6730 
C821 
6911 

6149 
6253 
6355 

6454 
6551 
6646 

6739 
6830 
6920 

6160 
6263 
6365 

6464 
6561 
6656 

6749 
6839 
6928 

6170 
6274 
6375 

6474 
6571 
6665 

6758 
6848 
6937 

6180 
6284 
C3S5 

6484 
6580 
6675 

6767 
6857 
6946 

6191 
6294 
6395 

6493 
6590 
6684 

6776 
6866 
6955 

6201 
6304 
6405 

6503 
6599 
6693 

6785 
6875 
6964 

6212 
6314 
6415 

6513 
6609 
6702 

6794 
6884 
6972 

6222 
6325 
6425 

6522 
6618 
6712 

6803 
6893 
6981 

12  3 
12  3 
12  3 

12  3 
1  2  3 
12  3 

12  3 

12  3 

12  3, 

4  6  6 
4  5  6 
4  6  6 

4  5  6 
4  6  6 
4  6  6 

4  6  6 

4  5  6 
4  4  5 

7  8  9 
7^9 
7  8  9 

7  8  9 
7  8  9 
7  7  8 

7  7  8 
7  7  8, 
6  7  8 

50 

6990 

6998 

7007 

7016 

7024 

7033 

7042 

7050 

7059 

7067 

12  3 

3  4  5 

6  7  8 

51 
52 
53 

54 

7076 
7160 
7243 

7324 

7084 
7168 
7251 

7332 

7093 
7177 
7259 

7340 

7101 
7185 
7267 

7348 

7110 
7193 

7275 

7356 

7118 
7202 
7284 

7364 

7126 
7210 
7292 

7372 

7135 
7218 
7300 

7380 

7143 
7226 
7308 

7388 

7152 
7235 
7316 

7396 

1  2  3 
12  3 
12  2 

12  2 

3  4  5 
3  4  5 
3  4  5 

3  4  6 

6  7  8 
6  7  7 
6  6  7 

6  6  7 

F 

0 

1 

S 

3 

4 

5 

6 

7 

8 

9 

1  2  2 

4  5  6 

7  8  9 

The  proportional  parts  are  stated  in  full  for  every  tenth  at  the  right-hand  side. 
Thb  logarithca  of  any  number  cf  four  significant  figures  can  be  read  directly  by  add* 


A] 


Table  A— Four  Place  Logarithms 


211 


IT 

0 

1 

2 

3 

4 

6 

6 

7 

8 

9 

12  3 

4  5  6 

7  8  9 

55 

56 

57 
58 
59 

7404 
7482 

7559 
7G34 
7709 

7412 
7490 

7566 
7642 
7716 

7419 
7497 

7574 
7649 
7723 

7427 
7505 

7582 
7657 
7731 

7435 
7513 

7589 
7664 
7738 

7443 
7520 

7597 

7672 
7745 

7451 
7628 

7604 
7679 
7752 

7459 
7636 

7612 
7686 
7760 

7466 
7543 

7619 
7694 
7767 

7474 
7651 

7627 
7701 

7774 

12  2 
1  2  2 

112 
112 

112 

3  4  5 
3  4  6 

3  4  5 
3  4  4 
3  4  4 

5  ()  7 

5  6  7 

6  6  7 
5  6  7 

5  6  7 

60 

7782 

7789 

7796 

7803 

7810 

7818 

7825 

7832 

7839 

7846 

112 

3  4  4 

5  6  6 

61 
62 
63 

64 
65 

66 

67 
68 
69 

7853 
7924 
7993 

8062 
8129 
8195 

8261 
8325 
8388 

7860 
7931 
8000 

8069 
8136 
8202 

8267 
8331 
8395 

7868 
7938 
8007 

8075 
8142 
8209 

8274 
8338 
8401 

7875 
7945 
8014 

8082 
8149 
8215 

8280 
8344 
8407 

7882 
7952 
8021 

8089 
8156 
8222 

8287 
8351 
8414 

7889 
7959 
8028 

8096 
8162 
8228 

8293 
8357 
8420 

7896 
7966 
8035 

8102 
8169 
8235 

8299 
8363 
8426 

7903 
7973 
8041 

8109 
8176 
8241 

8306 
8370 
8432 

7910 
7980 
8048 

8116 
8182 
8248 

8312 
8376 
8439 

7917 
7987 
8055 

8122 
8189 
8254 

8319 
8382 
8445 

112 
112 
112 

112 
112 
112 

112 
112 
112 

3  3  4 
3  §  4 
3  3  4 

3  3  4 
3  3  4 
3  3  4 

3  3  4 
3  3  4 
3  3  4 

6  6  6 
6  6  6 
5  5  6 

5  6  6 

6  6  6 
6  6  C 

5  6  6 
4  6  6 
4  5  0 

70 

8451 

8457 

8463 

8470 

8476 

8482 

8488 

8494 

8500 

8506 

112 

3  3  4 

4  5  6 

71 
72 
73 

74 
75 
76 

77 
78 
79 

8513 
8573 
8633 

8692 
8751 
8808 

8865 
8921 
8976 

8519 
8579 
8639 

8698 
8756 
8814 

8871 
8927 
8082 

8525 
8585 
8645 

8704 
8762 
8820 

8876 
8932 
8987 

8531 
8591 
8651 

8710 
8768 
8825 

8882 
8938 
8993 

8537 
8597 
8657 

8716 
8774 
8831 

8887 
8943 
8998 

8543 

8603 
8663 

8722 
8779 
8837 

8893 
8949 
9004 

8549 
8609 
8669 

8727 
8785 
8842 

8899 
8954 
9009 

8555 
8615 
8675 

8733 
8791 
8848 

8904 
8960 
9015 

8561 
8621 
8681 

8739 
8797 
8854 

8910 
8965 
9020 

8567 
8627 
8686 

8745 
8802 
8859 

8915 
8971 
9025 

112 
112 
112 

112 
112 
1  1  2 

112 
112 
112 

3  3  4 
3  3  4 
2  8  4 

2  3  4 
2  3  3 
2  3  3 

2  3  3 
2  3  3 
2  3  3 

4  6  6 
4  6  6 
4  6  6 

4  6  6 

4  6  5 
4  4  6. 

4  4  6 
4  4  5 
4  4  6 

80 

9031 

9036 

9042 

9047 

9053 

9058 

9063 

9069 

9074 

9079 

112 

2  3  3 

4  4  5 

81 
82 
83 

84 
85 
86 

87 
88 
89 

9085 
9138 
9191 

9243 
9294 
9345 

9395 
9445 
^94 

9090 
9143 
9196 

9248 
9299 
9350 

9400 
9450 
9499 

9096 
9149 
9201 

9253 
9304 
9355 

9405 
9455 
9504 

9101 
9154 
9206 

9258 
9309 
9360 

9410 

9460 
9509 

9106 
9159 
9212 

9263 
9315 
9365 

9415 
9465 
9513 

9112 
9165 
9217 

9269 
9320 
9370 

9420 
9469 
9518 

9117 
9170 
9222 

9274 
9325 
9375 

9425 
9474 
9523 

9122 

9175 
9227 

9279 
9330 
9380 

9430 
9479 
9528 

9128 
9180 
9232 

9284 
9335 
9385 

9435 
9484 
9533 

9133 
9186 
9238 

9289 
9340 
9390 

9440 
9489 
9538 

112 
112 
112 

112 
112 
112 

112 

Oil 
0  1  1 

2  3  3 
2  3  3 
2  3  3 

2  3  3 
2  3  3 
2  3  3 

2  3  3 

3  2  3 
2  2  3 

4  4  6 
4  4  5 
4  4  6 

4  4  6 
4  4  6 
4  4  6 

4  4  6 
3  4  4 

3  4  4 

90 

9542 

9547 

9552 

9557 

9562 

9506 

9571 

9576 

9581 

9586 

0  1  1 

2  2  3 

3  4  4 

91 
92 
93 

94 
95 

96 

97 
98 
99 

9590 
9638 
96^5 

9731 

9777 
9823 

9868 
9912 
9956 

9595 
9643 
9689 

9736 

9782 
9827 

9872 
9«)17 
9961 

960.) 
9647 
9694 

9741 
9786 
9832 

9877 
9921 
9965 

9605 
9652 
9699 

9745 
9791 
9836 

9881 
9926 
9969 

9609 
9657 
9703 

9750 
9795 
9841 

9886 
9930 
9974 

9614 
9661 
9708 

9754 
9800 
9845 

9890 
9934 
9978 

9619 
9666 
9713 

9759 
9805 
9850 

9894 
9939 
9983 

9624 
9671 
9717 

9763 

9809 
9854 

9899 
9943 
9987 

9628 
9675 
9722 

9768 
9814 
9859 

9903 
9948 
9991 

9633 
9680 
9727 

9773 

9818 
9863 

9908 
9952 
9996 

0  1  1 
0  1  1 

oil 
oil 

0  1  1 
0  1  1 

0  1  1 
0  1  1 

oil 

2  2  3 
2  2  3 
2  2  3 

2  2  3 
2  2  3 
2  2  3 

2  2  3 
2  2  3 
2  2  3 

3  4  4 
3  4  4 
3  4  4 

3  4  4 
3  4  4 
3  4  4 

3*4 
3  3  4 
3  3  4 

N 

0 

1 

2 

3 

4 

5 

6 

7 

8 

9 

12  3 

4  5  6 

7  8  9 

ing  the  proportional  part  corresponding  to  the  fourth  figure  to  the  tabular  number 
corresponding  to  the  first  three  figurest    There  may  be  au  error  uf  liu  the  last  place. 


212 


Table  B— Antilogarithms  to  Four  Places 


[B 


0 

1 

2 

3 

4 

5 

6 

7 

8 

9 

1 

2  3 

5 

6 

7  8  9 

.00 

1000 

1002 

1005 

1007 

1009 

1012 

1014 

1016 

1019 

1021 

0 

0  1 

1 

1 

2  2  2 

.01 
.02 
.03 

1023 
1047 
1072 

1026 
1050 
1074 

1028 
1052 
1076 

1030 
1054 
1079 

1033 
1057 
1081 

1035 
1059 
1084 

1038 
1062 
1086 

1040 
1054 
1089 

1042 
1067 
1091 

1045 
1069 
1094 

0 
0 
0 

0  1 
0  1 
0  1 

1 

1 
1 

2  2  2. 
2  2  2 
2^  2  2 

.04 
.05 

.06 

1096 
1122 
1148 

1099 
1125 
1151 

1102 
1127 
1153 

1104 
1130 
1156 

1107 
1132 
1159 

11Q9 
113^ 
1161 

1112 
1138 
1164 

1114 
1140 
1167 

1117 
1143 
1169 

1119 
1146 
1172 

0 
0 
0 

2 
2 
2 

2  2  2 
2  2  2 
2  2  2 

.07- 
.08 
.09 

1175 
1202 
1230 

1178 
1205 
1233 

1180 
1208 
1236 

1183 
1211 
1239 

1186 
1213 
1242 

1189 
1216 
1245 

1191 
1219 
1247 

1194 
1222 
1250 

1197 
1225 
1253 

1199 
1227 
1256 

0 
0 
0 

2 

2 
2 

2  2  2 
2  2  3 
2  2  3 

.10 

1259 

1262 

1265 

1268 

1271 

1274 

1276 

1279 

1282 

1285 

0 

1 

2 

2  2  3 

.11 
.12 
.13 

1288 
1318 
1349 

1291 
1321 
1352 

1294 
1324 
1355 

1297 
1221 
1358 

1300 
1330 
1361 

1303 
1334 
1365 

1306 
1337 
1368 

1309 
1340 
1371 

1312 
1343 
1374 

1315 
1346 
1377 

0 
0 
0 

2 
2 

2 

2 

2  2  3 
2  2  3 
2  3  3 

.14 
.15 

.16 

1380 
1413 
1445 

1384 
1416 
1449 

1387 
1419 
1452 

1390 
1422 
1455 

1393 
1426 
1459 

1396 
1429 
1462 

1400 
1432 
1466 

1403 
1435 
1469 

1406 
1439 
1472 

1409 
1442 
1476 

0 
0 
0 

2 
2 
2 

2 

2 
2 

2  3  .3 

2  3  3 
2^3 

.17 
.18 
.19 

1479 
1514 
1549 

1483 
1517 
1552 

1486 
1521 
1556 

1489 
1524 
1560 

1493 

1528 
1563 

1496 
1531 
1567 

1500 
1535 
1570 

1503 
1538 
1574 

1507 
1542 
1578 

1510 
1545 
1581 

0 
0 
0 

2 
2 
2 

2 
2 
2 

2  3  3 
2  3  3 
2  3  3 

.20 

1585 

1589 

1592 

1596 

1600 

1603 

1607 

1011 

1614 

1618 

0 

0 
0 
0 

2 

2 

3  3  3 

.21 
.22 
.23 

1622 
1660 
1698 

1626 
1663 
1702 

1629 
1667 
1706 

1633 
1671 
1710 

1631; 
1675 
1714 

1641 

1679 
1718 

1644 
1683 
1722 

1648 
1687 
1726 

1652 
1690 
1730 

1656 
1694 
1734 

2 

2 

2 

2 
2 

2 
2 

2 

3  3  3 
3  3  3 
3  3  3 

.24 
.25 

.26 

1738 
1778 
1820 

1742 

1782 
1824 

1746 
1786 
1828 

1750 
1791 
1832 

1754 
,1795 

'1837 

1758 
1799 
1841 

1762 
1803 
1845 

1766 
1807 
1849 

1770 
1811 
1854 

1774 

1816 
1858 

0 
0 
0 

2 

2 
2 

2 
2 
2 

2 
3 
3 

3  3  4 
3  3  4 
3  3  4 

.27 

.28 
.29 

1862 
1905 
1950 

1866 
1910 
1954 

1871 
1914 

1959 

1875 
1919 
1963 

1879 
1923 
19(>8 

1884 
1928 
1972 

1888 
1932 
1977 

1892 
1936 
1982 

1897 
1941 

1986 

1901 
1945 
1^)91 

0 
0 
0 

2 
2 

2 

2 
2 
2 

3 
3 
3 

3  3  4 
3  4  4 
3  4  4 

.30 

1995 

2000 

2004 

2009 

2014 

2018 

2023 

2028 

2032 

2037 

0 

2 
2 

2 

2 

2 
2 
2 

3 

3 
3 
3 

3  4  4 

.31 
.32 
.33 

2042 
2089 
2138 

2046 
2094 
2143 

2051 
2099 
2148 

205f) 
2104 
2153 

2061 
2109 
2158 

2065 
2113 
2163 

2070 
2118 
2108 

2075 
2123 
2173 

2080  2084 
2128  2133 
2178|  2183 

0 
0 
0 

3  4  4 
3  4  4 
3  4  4 

.34 
.35 

.36 

2188 
2239 
2291 

2193 
2244 
2296 

2198 
2249 
2301 

2203 
2254 
2307 

2208 
2259 
2312 

2213 
2205 
2317 

2218 
2270 
2323 

2223 

2275 
2328 

2228  2234 
2280  2280 
2333  2339 

1  2 

1  2 

1  2 

2 
2 
2 

3 
3 
3 

3 
3 
3 

4  4  5 
4.4  5 
4  4  5 

.37 
.38 
.39 

2344 
2399 
2455 

2350 
2404 
2460 

2355 
2410 
2466 

2360 
2415 

2472 

2366 
2421 

2477 

2371 
2427 

2483 

2377 
2432 
2489 

2382 
2438 
2495 

2388 
2443 
2500 

2393 
2449 

2506 

1  2 

1  2 
1  2 

2 

2 

2 

3 
3 
3 

3 
3 
3 

4  4  5 
4  5  5 
4  5  5 

.40 

2512 

2518 

2523 

2529 

2535 

2541 

2547 

2553 

2559 

2564 

1  2 

2 

3 

4 

4  5  5 

.41 
.42 
.43 

2570 
2630 
2692 

2576 
2636 
2698 

2582 
2642 
2704 

2588 
2649 
2710 

2594 
2655 
2716 

2600 
2661 
2723 

2606 
2067 
2729 

2612 

2673 
2735 

2618 
2679 
2742 

2624 
2685 
2748 

1  2 
1  2 
1  2 

2 
2 
2 

3 

3 
3 

4 
4 
4 

4  5  6. 
4  5  6 
4  6  6 

.44 
.45 

.46 

2754 
2818 
2884 

2761 
2825 
2891 

2767 
2831 
2897 

2773 
2838 
2904 

2780 
2844 
2911 

2786 
2851 
2917 

2793 
2858 
2924 

2799 
2864 
2931 

2805 
2871 
2938 

2812 
2877 
2944 

1  2 
1  2 
1  2 

3 
3 
3 

3 
3 
3 

4 
4 
4 

4  5  6 

5  5  6 
5  5  6 

.47 

.48 
.49 

2951 
3020 
3090 

2958 
3027 
3097 

2965 

3o:^ 

3105 

2972 
3041 
3112 

2979 
3048 
3119 

2985 
3055 
3126 

2992 
3062 
3133 

2999 
3069 
3141 

3006 
3Q76 
3148 

3013 
3083 
3155 

1  2 

1  2 

1  2 

3 

3 
3 

3 
3 
4 

4 
4 
4 

5  6  6 
5  6  6 
5  6  6 

'    <.   r   t   A 


B] 

Table  B- 

—  Afitilogarithms  to  Four  Places 

il3 

0 

1 

2 

3 

4 

5 

6 

7 

8 

9 

1.2  3 

4  5  6 

7  8  9 

.50 

3162 

3170 

3177 

,3184 

3192 

3199 

3206 

3214 

3221 

3228 

1  1  2 

3  4  4 

5  6  7 

.51 
.52 
.53 

.54 
.65 
.56 

.57 
.58 
.59 

3236 
3311 
3388 

3467 
3548 
3631 

3715 

3802 
3890 

3243 
3319 
3396 

3475 
3556 
3639 

3724 
3811 
3899 

3251 
3327 
3404 

3483 
3565 
3648 

3733 
3819 
3908 

3258 
3334 
3412 

3491 
3573 
3656 

3741 

3828 
3917 

3266 
3342 
3420 

3499 
3581 
3664 

3750 
3837 
3926 

3273 
3350 
3428 

3508 
3589 
3673 

3758 
3846 
3936 

3281 
3357 
3436 

3516 
3597 
3681 

3767 
3855 
3945 

3289 
3365 
3443 

3524 
3606 
3690 

3776 
3864 
3954 

3296 
3373 
3451 

3532 
3614 
3698 

3784 
3873 
3963 

3304 
3381 
3459 

3540 
3622 
3707 

3793 

3882 
3972 

112 
1  1  2 
12  2 

12  2 
12  2 
12  2 

12  3 
12  3 
12  3 

3  4  4 
3  4  5 
3  4  5 

3  4^5' 
3  4  5 
3  4  5 

3  4  5 

3  4  5 

4  5  5 

5  6  7 

5  6  7 

6  6  7 

6  6  7 
6  7  7 
6  7  8 

6  7  8^ 
6  7  8 
6  7  8 

.60 

3981 

3990 

3999 

4009 

4018 

4027 

4036 

4046 

4055 

4064 

12  3 

4  5  6 

7  8  8 

.61 
.62 
.63 

.64 
.65 

.66 

.67 
.68 
.69 

4074 
4169 
4266 

4365 
4467 
4571 

4677 
4786 
4898 

4083 
4J7^ 
4276 

4375 
4477 
4581 

4688 
4797 
4909 

4093 
4188 
4285 

4385 

4487 
4592 

4699 
4808 
4920 

4102 
4198 
4295 

4395 
4498 
4603 

4710 
4819 
4932 

4111 
4207 
4305 

4406 
4508 
4613 

4721 

4831 
4943 

4121 
4217 
4315 

4416 
4519 
4624 

4732 
4842 
4955 

4130 
4227 
4325 

4426 
4529 
4634 

4742 
4853 
4966 

4140 
4236 
4335 

4436 
4539 
4645 

4753 
4864 
4977 

4150 
4246 
4345 

4446 
4550 
4656 

4764 
4875 
4989 

4159 
4256 
4355 

4457 
4560 
4667 

4775 
4887 
5000 

12  3 
12  3 
12  3 

12  3 
1  2  3 
12  3 

12  3 
1  2  3 
1  2  3 

4  5  6 
4  5  6 
4  5  6 

4  5  6 
4  5  6 
4  5  6 

4  5  7 

5  6  7 
5  6  7 

7  8  9 
7  8  9 
7  8  9 

7  8  9 

7  8  9 

7  910 

8  910 
8  910 
8  9  10 

.70 

5012 

5023 

5035 

5047 

5058 

5070 

5082 

5093 

5105 

5117 

12  3 

5  6  7 

8  910 

.71 
.72 
.73 

.74 
.76 

.76 

.77 
.78 
.79 

5129 
5248 
5370 

5495 
5623 
5754 

5888 
6026 
6166 

5140 
5260 
5383 

5508 
5636 
5708 

5902 
6039 
6180 

5152 
5272 
5395 

5521 
5649 
5781 

5916 
6053 
6194 

5164 
5284 
5408 

5534 
5662 
5794 

5929 
6067 
6209 

5176 
5297 
5420 

5546 
5675 
5808 

5943 
6081 
6223 

5188 
5309 
5433 

5559 
5689 
5821 

5957 
6095 
6237 

5200 
5321 
5445 

5572 
5702 
5834 

5970 
0109 
6252 

5212 
5333 
5458 

5585 
5715 
5848 

5984 
6124 
6266 

5224 
5346 
5470 

5598 
5728 
5861 

5998 
6138 
6281 

5236 
5358 
5483 

5610 
5741 
6875 

6012 
6152 
6295 

12  4 

1  2  4 

13  4 

13  4 
13  4 
13  4 

13  4 
13  4 
13  4 

5  6  7 
5  6  7 
5  6  7 

5  6  8 
5  7  8 
5  7  8 

5  7  8 

6  7  8 
6  7  9 

81011 
910  11 
91011 

91012 
91112 
91112 

10  11 12 
10  11 13 
10  11 13 

.80 

6310 

6324 

6339 

6353 

6368 

6383 

6397 

6412 

6427 

6442 

13  4 

6  7  9 

10  12  13 

.81 
.82 
.83 

.84 
.85 

.86 

.87 
.88 
.89 

.90 

6457 
6607 
6761 

6918 
7079 
7244 

7413 

7586 
7762 

6471 
6622 
6776 

6934 
7096 
7261 

7430 
7603 
7780 

6486 
6v037 
6792 

6950 
7112 

7278 

7447 
7621 
7798 

6501 
6653 
6808 

6966 
7129 
7295 

7464 
7638 
7816 

6516 

6668 
6823 

6982 
7145 
7311 

7482 
7656 
7834 

6531 

6683 
6839 

69G8 
7161 

7328 

7499 
7674 
7852 

6546 
6699 
6855 

7015 
7178 
7345 

7516 
7691 

7870 

6561 
6714 
6871 

7031 
7194 
7362 

7534 
7709 
7889 

6577 
6730 
6887 

7047 
7211 
7379 

7551 

7727 
7907 

6592 
6745 
6902 

7063 
7228 
7396 

7568 
7745 
7925 

2  3  5 
2  3  5 
2  3  5 

2  3  5 
2  3  5 
2  3  5 

2  4  5 
2  4  5 
2  4  6 

6  8  9 
6  8  9 

6  8  9 

7  810 
7  810 
7  810 

7  910 
7  911 
7  911 

11 12  14 

11  12  14 

11 13  14 

11 13  15 

12  13  15 
12  14  15 

12  14  16 

12  14  16 

13  15  16 

7943 

7962 

7980 

7998 

8017 

8035 

8054 

8072 

8091 

8110 

2  4  6 

7  911 

13  15  17 

.91 
.92 
.93 

.94 
.95 

.96 

.97 

.98 
.99 

8128 
8318 
8511 

8710 
8913 
9120 

9333 
9550 
9772 

8147 
8337 
8531 

8730 
8933 
9141 

9354 
9572 
9795 

8166 
8356 
8551 

8750 
8954 
9162 

9376 
9594 
9817 

8185 
8375 
8570 

8770 
8974 
9183 

9397 
9616 
9840 

8204 
8395 
8590 

8790 
8995 
9204 

9419 
96;i8 
9863 

8222 
8414 
8610 

8810 
iK)16 
9226 

9441 
9661 
9886 

8241 
8433 
8630 

8831 
9036 
9247 

9462 
9(>83 
9908 

8260 
8453 
8650 

8851 
9057 
9208 

9484 
9705 
9931 

8279 
8472 
8670 

8872 
9078 
9290 

9506 
9727 
9954 

8299 
849*2 
8690 

8892 
9099 
9311 

9528 
9750 
9977 

2  4  6 
2  4  6 
2  4  6 

2  4  6 
2  4  6 
2  4  6 

2  4  6 
2  4  7 

2  5  7 

8  911 
8  10  12 
81012 

8  1012 
8  10  12 
91113 

91113 
91113 
91114 

13  15  17 

14  15  17 
14  16  18 

14  16  18 

15  17  19 
15  17  19 

15  17  19 

16  18  20 
1618J^ 

INDEX 


/ 


Accuracy 4 

Acidimetry 107,  118 

Acids,  degree  of  ionization  of     .     .       25 
determination  of  available  hy- 

•    drogen-ion  in 118 

standard  solutions  of  .     .     .     .     11 1 

titration  of 108 

Adsorption 23,  31 

AflSnity  constant 26 

Afterflow,  error  from 45 

Alkalimeter,  Mohr's   .     .     .     .     71,  72 

Alkalimetry 107,  116,  120 

Alkali  solutions,  standard  .  .  .  in 
Aluminum,  determination  of  .  .  63 
Ammonium  thiocyanate,  standard 

solutions  of 166 

Ampere,  definition  of 92 

Analyzed  chemicals 3 

Antilogarithms 212 

Antimony,    determination    of    in 

stibnite 155 

Apparatus,  list  of  for  quantitative 

work 209 

Arsenic,  removal  of  from  copper  .     163 

Arsenious  oxide,  primary  standard     152 

Asbestos,  preparation  of  for  filters      34 

use  of  in  filtration   .     .     .     .     34,  35 

Ashless  filter  papers 31 

Atomic  weights,  table  of        Back  cover 

sheet 

Baking  powder,  determination  of 

carbon  dioxide  in      .     ,     .     73,  74 

Balance,  analytical 7 

adjustment  of 9 

conditions  to  be  fulfilled  by  .     .  8 

exercises  with 53 

location  of 8 

relative  length  of  arms  of .     .     .  18 

rest-point  of 11 

sensitiveness  of 13 

use  and  care  of 9 

Barium  sulphate,  properties  of  .     64,  65 


Bases,  degree  of  ionization  of    .    .  25 

standard  solutions  of    .     .     .     .  iii 

titration  of 108 

Bumping       42 

Buoyancy,  correction  for    .     .     .  19 

Burettes 43 

calibration  of 48 

cleaning  of 45 

reading  of 45 

Burning  filter  papers 39 

Calcium,  determination  of  in  lime- 
stone       75.  137 

oxalate,  properties  of  ...     .  78 

Calibration  of  measuring  vessels  .  48 

of  weights 15 

Carbon  dioxide,  determination  of 

in  limestone 71 

determination  of  in  baking  pow- 
ders       73,  74 

Chemical  equilibrium      ....  25 

Chemical  factors 186 

Chlorine,  determination  of   55,  59,  168 
Chrome  iron  ore,  determination  of 

chromitmi  in 157 

Chromic  oxide,  ignition  of    ...  63 

Chromium,  determination  of      63,  157 
Cleaning  solution,  preparation  and 

use  of 208 

Colloidal  precipitates 22 

Common-ion  effect 27 

Contamination  of  precipitates  .     .  23 

Contat-Gockel  valve    .     .     .     138,  139 

Copper,  determination  of     .     .  89,  162 
Counterpoise,  use  of  in  weighing 

bulky  objects       .     .     .     .     21,  74 

Crucibles,  materials  of    ...     .  40 

Current  density 96 

Current,  production  of  for  electro- 
analysis      92 

Decantation 32 

Decimal,  number  of  places  to  re- 
port    6 


21S 


2l6 


INDEX 


Deposition  voltages  of  elements, 

table  of 94 

Desiccators 39 

Dichromate  processes      .     .     .     .  125 

solutions,  standard 125 

Dichromate-sulphuric  acid  clean- 
ing solution 208 

Digestion  of  precipitates .     .     .     ^  22 

Distilled  water,  testing  of     .     .     .  3 

Double  precipitation 24 

Drainage,  error  from 45 

Drying  ovens 38 

Electro-analysis 89 

Electrode  potentials,  table  of  .  .  94 
Electrodes,  material  and  form  of  .  97 
Electrolytes,  influence  of  composi- 
tion of  upon  electro-analysis  .  96 
Electrolytic  separations  ....  94 
Electrolytic  solution  tension  .  .  93 
End-point  in  titration,  determina- 
tion of 102 

Equilibrium,  chemical     ....  25 

Equilibrium  constant 26 

Evaporation  of  liquids    ....  42 

Factor,  chemical 186 

normality       190 

Faraday's  laws 92 

Ferric  alum,  indicator  .  .  r66,  167 
Ferric  oxide,  ignition  of  .  .  .  61,  63 
Ferrous      ammonium      sulphate, 

standard  solutions  of    .     .     .  125 
Fertilizers,  determination  of  nitro- 
gen in 120,  122 

Filters,  selection  and  use  of .     .     .  31 

Filtrates,  testing  of 33 

Filtration 31 

Fine-grained  precipitates,  enlarge- 
ment of  the  particles  of      .     .  22 

Flasks,  volumetric 44 

calibration  of 48 

Flocculation  of  colloids    ....  22 

Funnels,  selection  of 32 

Fusions,  removal  of  from  crucibles  85 

Glassware 3 

Gooch  filters,  preparation  and  use 

of 34,  35 

sources  of  error  with    ....  36 


Graduated  cylinders 44 

Gravimetric  analysis i,  S3 

Halogens,  determination  of .     .     .  59 
Hematite,  determination  of  iron 

in 133 

Hydrochloric  acid,  standard  solu- 
tions of  112 

Ignition  of  precipitates    ....  38 

Indicators       ....     102,  107,  124 

general  theory  of 103 

sensitiveness  of  in  alkaUmetry 

and  acidimetry     .     .     .     108,  109 

Indirect  methods  of  analysis    .     .  191 
Insoluble  matter,  determination  of 

in  limestone 77 

International     atomic      weights, 

table  of      .     .     .     Back  cover  sheet 

Iodine  standard  solutions  of      148,  151 

lodometric  processes 148 

Ionization,  degree  of 25 

repression  of 27 

Ions,  complex 25 

composition  of  the 25 

Iron,  determination  of     .  60,  129,  133 
oxidation  of  ferrous  to  ferric 

61,  62,  127,  133 
reduction  of  ferric  to  ferrous 

63,  127,  135 

Iron  ores, decomposition  of  129, 134,  135 

Iron  wire,  primary  standard     .     .  126 

Jones  reductor,  assembly  and  use 

of 142 

Kjeldahl,    determination   of    pro- 
tein nitrogen   ......  120 

Labels 4 

Laboratory  Records 5 

Lead,  determination  of  in  an  ore    .  159 

Limestone,  analysis  of    71,  75,  77,  137 
determination  of  carbon  dioxide 

in 71 

Limits  of   error  in   experimental 

work 5 

Liquids,  evaporation  of   ...     .  42 

transference  of 42 

volumetric  measurement  of  .     .  43 


INDEX 


217 


Liter,  Mohr's 47 

normal 47 

true  .  ' .' 48 

Logarithms 210 

Magnesium,  determination  of  in 

limestone 75.  79 

Magnesium      ammonium      phos- 
phate, ignition  of     ...     .  80 
Manganese,  determination  of  in  an 

ore 144 

Manganese  ores,  decomposition  of  146 

Methyl  orange  solution  .     .     .     .  no 

Molybdic  acid,  recovery  of  .     .     .  207 

Neutral  solution,  definition  of  .     .  107 

Neutralization  methods  ....  107 

Nickel ,  electrolytic  determination  of  91 
Nitrogen,  Kjeldahl  determination 

of 120 

Normality  factor,  definition  of      .190 

Normal  solutions 100 

Normal  System  of  Reagents,  ad- 
vantages of 205 

Ohm,  definition  of 92 

Ohm's  law 92 

Ovens,  drying 38 

Overvoltage 94 

Oxidation  and  reduction  methods  124 
Oxidizing  agents,   standard  solu- 
tions of  124 

Parallax,  error  from 45 

Permanganate  processes      .     .     .  130 
Permanganate  solutions,  standard  131 
Phenolphthalein  solution     .     .     .  no 
Phosphoric  anhydride,  determina- 
tion of        81 

Phosphorus,  determination  of  in 

steel 140 

Pipettes,  transfer 44 

calibration  of 48 

Platinum  ware,  defects  of  modern  41 

specifications  for 41 

use  and  care  of    .     .     .     .40,  41,  89 
Polarity  of  terminals,  determina- 
tion of '    .  89 

Polarization 94 

Policeman,  definition  of  ...     .  ^:s 

Potash,  determination  of     .     .     .  68 


Potassium     bitartrate,      primary 

standard 115 

dichromate,  standard  solutions 

of 125 

ferricyanide,  indicator  .     .     125,  126 
iodate,  primary  standard  .     .     .  153 
permanganate,    standard    solu- 
tions of  131 

thiocyanate,  standard  solutions 

of 166 

Precipitates,  colloidal      ....  22 

contamination  of 23 

digestion  of 22 

drying  of 38 

enlargement  of  the  particles  of  .  22 

fine-grained 22 

flocculation  of 22 

for  use  in  gravimetric  analysis  .  21 

ignition  of 38,  39 

purification  of 24 

washing  of     .     .     .      21,  31,  33,  36 

Precipitation 21 

theory  of 24 

voltunetric  methods  of     .     .     .  165 

Problems 192 

Pyrolusite,  Mn02  value  of    ...  138 

Questions 170 

Reactions  of  volumetric  analysis  .  loi 

Reagents,  analyzed 3 

preparation  of 205 

quality  of 3 

testing  of 3 

Records 5 

Reducing  agents,  standard  solu- 
tions of  124 

Reduction,  methods  of  oxidation 

and 124 

Reductor,  Jones 142 

Rest-point  of  balance,  determina- 
tion of II 

Reversible  reactions 24 

Salts,  degree  of  ionization  of    .     .  25 
Samples,  Preparation  of,  for  An- 
alysis          so,  208 

Saturated  solution,  definition  of  .  29 

Sensitiveness,  of  balance     ...  13 

of  indicators,  table  of  the  .     .     .  109 
Siderite,  determination  of  iron  in  . 


2l8 


INDEX 


Silica,  determination  of  ...  .  85 
Silicates,  determination  of  silica  in 

refractory - .     .      85 

Silicic  acid,  dehydration  of  .  .  .  87 
Silver,  determination  of  .  .  .  59,  165 
Silver,  primary  standard  .  .  .  167 
Silver  chloride,  properties  of  .  57,  58 
Silver  ion,  properties  of  ...  .  58 
Silver  nitrate,  standard  solutions 

of 166 

Soda,  alkaline  value  of    .     .     .     .     116 
Sodium  carbonate,  primary  stand- 
ard          113,  114 

Sodium  chloride,  determination  of 

chlorine  in 55,  168 

primary  standard 166 

purification  of 167 

Sodium  hydroxide,  standard  solu- 
tions of 112 

Sodium  oxalate,  primary  standard     132 
Sodium  thiosulphate,  standard  so- 
lutions of      151,  153,  161,  162,  164 
Solubility,  effect  of  size  of  particles 

on 22 

Solubility  product 28 

Solution  tension 28 

electrolytic 93 

Solution  of  iron  ores  .  .129,  134,  135 
Solution  of  manganese  ores  .  .  146 
Standardization,  definition  of  .  .  99 
of  hydrochloric  acid  .  .  .  .  112 
of  sodium  hydroxide  solution  .  112 
of  dichromate  solution  .  .  .  125 
of  ferrous  ammonium  sulphate 

solution 125 

of  iodine  solution 151 

of  permanganate  solution  .  .  131 
of  silver  nitrate  solution  .  .  .  166 
of  sodium  thiosulphate  solution 

151,  161,  162 
of  thiocyanate  solution     .     .     .     166 
Standard  solution,  definition  of     .       99 
Starch,  indicator    .     .     .     .      150,  151 
Stibnite,    determination   of   anti- 
mony in 155 

Suction,  use  of 32,  34,  35 

Sulphur,  determination  of   .     .    64,  66 


Temperature,   correction  for  dif- 
ferences in 46 

Tension,  solution 28 

electrolytic  solution     ....  93 

Testing  for  complete  precipitation  33 

of  washings 33 

Thiocyanate  solutions,  standard  .  166 

Titration,  definition  of     ...     .  2 

Transference  of  liquids    ....  42 

Transfer  pipettes 44 

Triangles 41,  42 

Utilization  of  time 4 

Vacuum,  use  of 32,  34,  35 

Valve,  Contat-Gockel     .     .      138,  139 

Volt,  definition  of 92 

Volume,  units  of 47 

Volumetric  analysis,  general  dis- 
cussion    99 

fundamental  reactions  of .     .     .  loi 
neutralization  methods  of     .     .  107 
oxidation  and  reduction  meth- 
ods of 124 

precipitation  methods  of  .     .     .  165 

Volumetric  apparatus     ....  43 

calibration   of 48 

cleaning  of 45 

sources  of  error  in  the  use  of  .     .  44 
Volumetric    System,    Advantages 

of          43.  105 

Volumetric    Work,    General    Re- 
marks      105 

Wash  bottles 33 

Washing  of  precipitates  .     .  22,  31,  33 

theory  of 36 

Washings,  testing  of 33 

Water,  ionization  of 107 

Weighing 7,  53 

Umits  of  error  in 11 

methods  of 12 

summary  of 20 

Weights,  calibration  of    ...     .  15 

use  and  care  of 10 

Zimmermann-Reinhardt  solution  .  134 


«. . 

5  1 

16 

6f 

->    5 

^'^ 

THIS  BOOK  IS  DUE  ON  THE  LAST  DATE 
STAMPED  BELOW 


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DEC  5    1967 
DEC  5    R^ 


LIBRARY,  UNIVERSITY  OF  CALIFORNIA,  DAVIS 

Book  Slip-50m-12,'64(F77284)458 


361537 

QDlOl 

Smith,    G.M. 

S6 

A  course  of  in- 

1921 

struction  in 
quantitative 

Aliimiii 

96.0 

Antimc 

chemical  anal- 

144.3 

Argon 

ysis. 

20.2 

Arsenic 

58.68 

Barium 

222.4 

Bismut 

14.008 

Boron 

190.9 

Bromin 

16.000 

Cadmiu 

106.7 

Caesium 

31.04 

Calciun 

195-2 

Carbon 

39-iP 

Cerium 

140.9 

Chlorin 
Chromi 

LIBRARY 
UNIVERSITY  OF  CALIFORNIA 

226.0 
102.9 

Cobalt 

PnliimK 

DAVIS 

85.45 

i^oiumDiuiu 
Copper 

Cu            63.57 

Samarium 

Sa 

101.7 
150.4 

Dysprosium 

Dy         162.5 

Scandium 

Sc 

45-1 

Erbium 

Er          167.7 

Selenium 

Se 

79.2 

Europium 

Eu         152.0 

SiUcon 

Si 

28.3 

Fluorine 

F             19,0 

Silver 

Ag 

107.88 

Gadolinium 

Gd         157.3 

Sodium 

Na 

23.00 

Gallium 

Ga           70.1 

Strontium 

Sr 

87.63 

Germanium 

Ge           72.5 

Sulphur 

S 

32.06 

Glucinum 

Gl              9.1 

Tantalum  \ 

Ta 

181. 5 

Gold 

Au         197.2 

Tellurium 

Te 

127.5 

HeUum 

He             4.00 

Terbium 

Tb 

159.2 

Holmium 

Ho         163.5 

Thallium 

Tl 

204.0 

Hydrogen 

H              1.008 

Thorium 

Th 

232.15 

Indium 

In          1 14.8 

ThuUum 

Tm 

168.5 

Iodine 

I            126.92 

Tin 

Sn 

118.7 

Iridium 

•    Ir           1931 

Titanium 

Ti 

48.1 

Iron 

Fe            55-84 

Tungsten 

W 

184.0 

Krypton 

Kr           82.92 

Uranium 

U 

238.2 

Lanthanum 

La          139.0 

Vanadium 

V 

51.0 

Lead 

Pb         207.20 

Xenon 

Xe 

130.2 

Lithium 

Li              6.94 

Ytterbium 

Yb 

173.5 

Lutetium 

Lu          175.0 

Yttrium 

Yt 

89-33 

Magnesium 

Mg          24.32 

Zinc 

Zn 

65.37 

Manganese 

Mn          54.93 

Zirconium 

Zr 

90.6 

Mercury 

r 

Hg         200.6 

